D 


UC-NRLF 


277 


LABORATORY  EXERCISE S 

TO  ACCOMPANY 

STORER  AND  LINGSM'S  CHEMISTRY 


LIBRARY 

OF  THE 

UNIVERSITY  OF  CALIFORNIA 


GIETT 


i 


LABORATORY   EXERCISES 


IN 


GENERAL     CHEMISTRY 


COMPILED  FROM  VARIOUS  SOURCES 
BY 

G.    W.    SHAW,   A.M. 

FORMERLY  PROFESSOR  OF  CHEMISTRY  AT   OREGON  STATE 
AGRICULTURAL    COLLEGE 


FOR  USE    IN   CONNECTION  WITH 
STOKER  AND    LINDSAY'S  MANUAL   OF   CHEMISTRY 


\  S  rf  A  f? 

OK   THE 


UNIVERSITY 

OF 

£lUFOR]^ 


NEW  YORK  •:•  CINCINNATI  •:•  CHICAGO 

AMERICAN    BOOK    COMPANY 


COPYRIGHT,  1901,  BY 
CK  W.  SHAW. 

ENTERED  AT  STATIONERS'  HALL,  LONDON. 

SHAW.      LAB.    EX. 
W.  P.  i 


PREFACE. 

THIS  pamphlet  claims  to  present  nothing  that  is  new  and 
untried.  Beginners  in  chemistry  need  such  experiments  as 
give  certain  and  well-defined  results.  The  pamphlet  has  been 
prepared  to  use  in  the  laboratory  in  connection  with  Storer  & 
Lindsay's  "  Manual  of  Chemistry  "  in  the  class  room,  hence  the 
greater  number  of  the  experiments  have  been  drawn  from  that 
work,  yet  many  from  other  sources  have  been  included.  A 
large  number  of  manuals  have  been  examined  in  the  prepara- 
tion of  the  pamphlet,  but  the  compiler  is  particularly  indebted 
to  those  of  Williams  and  Remsen. 

The  raison  d'&tre  for  the  book  is  simply  that  of  facilitating 
work  in  the  laboratory  by  having  before  the  student  only  such 
matter  as  is  necessary  for  the  work  in  hand.  It  has  been  the 
endeavor  of  the  compiler  to  include  only  such  experiments  as 
are  truly  instructive  and  well  illustrate  the  subject,  and  which 
lead  up  easily  and  naturally  to  a  knowledge  of  the  science. 
The  book  is  to  be  used  in  a  laboratory  in  which  an  instructor 
is  always  present  with  the  class  to  furnish  such  individual 
direction  as  may  be  necessary. 

Most  of  the  experiments  should  be  performed  on  the  lecture 
table  before  they  are  attempted  by  the  student. 

Any  text-book  may  be  consulted  outside  the  laboratory ;  in 

the  laboratory  itself  no  other  book  than  this  is  allowed.    In 

3 

J 07459 


4  PREFACE. 

some  cases  a  sketch  of  the  necessary  apparatus,  or  an  example 
of  the  apparatus  itself,  is  placed  in  the  laboratory,  where  it 
can  be  seen  by  the  students. 

The  last  twenty  minutes  of  the  laboratory  period  should  be 
taken  to  question  the  class  rapidly  on  the  work  of  the  day. 
The  books  should  be  inspected  after  each  experiment  has  been 
written  out. 


OF  THE 

UNIVERSITY 


TO  THE   STUDENT. 

1.  Unless  special  directions  to  the  contrary  are  given  no 
text-book    will    be    allowed    in   the    laboratory   except   this 
pamphlet. 

2.  Provide  yourself  with  a  working  apron  to  protect  your 
clothing ;  also  with  towel  or  cloth  to  clean  the  desk. 

3.  Neatness  is  essential  to  success  in  chemical  work.     Both 
the  desk  and  apparatus  must  be  kept  deem  at  all  times. 

4.  Pupils  are  held  responsible  for  apparatus,  and   must 
replace  anything  lost  or  broken. 

5.  Never  mix  chemicals  or  reagents  except  as  directed. 

6.  In  experimenting  have  all  apparatus  neatly  arranged. 
Have  every  stopper  and  connection  fit  tightly. 

7.  Have  flasks   and  tubes   perfectly  dry  on  the   outside 
before  applying  heat. 

8.  Keagents  for  general  use  must  not  be  taken  to  the  stu- 
dent's desk,  but  used   at  the   side   table.      Any  excess  of  a 
reagent  must  not  be  poured  back  from  a  test  tube  or  beaker  into 
the  reagent  bottle,  nor  should  a  .stirring  rod  be  dipped  into  a 
reagent  bottle. 

9.  In  experimenting  follow  directions  as  closely  as  pos- 
sible.     Kead  an  experiment  through  before  performing  any 
part  of  it.      Ask   an   explanation   of   anything  you   do  not 
understand. 

10.  Observe  very  carefully  everything  that  takes  place ; 
and  endeavor  to  distinguish  essential  from  non-essential  phe- 
nomena; express  in  writing  the  results  of  your  observation 
and  the  conclusions  as  to  facts  taught  by  the  experiment. 

6 


6  TO  THE   STUDENT. 

11.  A  student's   standing   is   largely  determined  by  the 
quality  of  his   laboratory  work  and  notebooks.       The   notes 
must  be  written  clearly  and  distinctly.     Neatness  will  be  in- 
sisted upon.      Do  not   crowd  the  notes,  but  leave  room  for 
remarks,  or   corrections.      In   general   answer  the  following 
questions :  — 

(a)  What  is  the  object  of  the  experiment  ? 

(b)  What  materials  did  you  use  ? 

(c)  What  apparatus  did  you  use  ? 

(d)  What  did  you  do  ? 

(e)  What  did  you  observe  ? 

(/)  What  are  your  conclusions  ? 

12.  At  the  end  of  the  term  no  allowance  is  made  for  appa- 
ratus that  is  not  clean  and  in  proper  condition  to  be  served 
out  to  other  students. 

I  have  read  the  above  directions  carefully,  and  endeavored 
to  understand  them. 

Signed  


LABORATORY   EXERCISES 


EXERCISE  I.     PRELIMINARY. 

§§  1  to  5.1 

1.  Ascertain  the  number  of  your  desk. 

2.  Obtain  key  to  desk  at  the  supply  room. 

3.  Examine  and  take  inventory  of  contents  of  desk. 

4.  Take  receipt  to  the  supply  room. 

Experiment  1.  Measure  in  a  graduate  10  cc.  of  water,  then  pour  it 
into  a  test  tube.  Note  what  proportion  of  the  tube  is  filled.  Pour 
out  the  water,  then  pour  into  the  tube  as  near  the  same  quantity  of 
water  as  possible,  estimating  by  the  eye  alone.  Verify  your  estimate 
by  measuring  in  a  graduate.  Repeat  till  you  can  estimate  closely. 

Hereafter,  unless  great  accuracy  is  required,  estimate  volumes  without 
measuring. 

Supplementary.  Learn  the  metric  table  for  length,  weight,  and  vol- 
ume. How  many  cubic  centimeters  in  a  liter  ?  What  is  the  weight  in 
grams  of  a  liter  of  water?  Mercury  weighs  13,6  times  as  much  as 
water.  What  will  5|  cc.  of  mercury  weigh  ?  A  rectangular  block  of 
marble  is  3x4x2  cm.,  and  has  a  sp.  gr.  2.5.  What  does  it  weigh 
in  grams? 

2.  (a)  Try  to  dissolve  a  small  quantity  of  sulphur  in  carbon  bisul- 
phide. Try  to  pick  up  small  pieces  of  iron  with  a  magnet.  Try  the 
effect  of  dilute  hydrochloric  acid  on  sulphur.  Is  any  gas  given  off  ? 
Try  the  same  acid  on  iron  particles. 

(&)  Mix  thoroughly  small  quantities  of  sulphur  and  iron.  Is  a  new 
substance  formed?  Is  the  sulphur  still  sulphur?  Is  the  iron  still 
iron? 

Weigh  out  3  g.  of  sulphur  and  6  g.  of  iron ;  mix  thoroughly,  and 
put  it  into  a  tube  of  hard  glass  closed  at  one  end.  Hold  the  tube  by 

1  Lesson  in  the  "  Manual  of  Chemistry  "  to  precede  the  exercise. 

7 


8  LABORATORY   EXERCISES. 

the  open  end  by  means  of  a  strip  of  paper  folded  several  times,  and 
heat  the  mixture  over  a  burner.  When  no  further  change  takes  place 
allow  the  tube  to  cool ;  break  it,  and  examine  the  contents  for  sulphur 
and  for  iron. 

3.  Heat  a  small  piece  of  lead  foil  on  the  lid  of  a  porcelain  crucible 
as  long  as  any  change  occurs.     Warm  slowly,  as  sudden  heat  will 
cause  the  porcelain  to  break.     While  this  is  going  on,  perform  the 
following :  — 

4.  Fit  to  any  small  flask  or  bottle  a  perforated  cork  to  which  has 
been  adapted  a  short  piece  of  glass  tubing.     Over  the  end  of  this  glass 
tube  slip  a  short  piece  of  caoutchouc  tubing.     Suck  part  of  the  air  out 
of  the  flask,  and  then  nip  the  caoutchouc  tubing  with  the  thumb  and 
finger  so  that  no  air  shall  reenter.     Immerse  the  neck  of  the  flask  in  a 
basin  of  water  and  release  the  tubing.     Explain  all  results. 

5.  Adapt  the  throat  of  a  funnel  to  a  perforated  cork  which  fits 
tightly  into  the  neck  of  a  bottle  or  flask,  and  then  fill  the  funnel  with 
water.     Why  does  the  water  not  enter  the  bottle  ? 

Loosen  the  cork  so  that  the  mouth   of  the  bottle  is  not  tightly 
closed.     Explain  the  action. 

6.  Burn  some  magnesium  wire,  observe  carefully  the  result  of  the 
combustion,  and  describe  the  product. 

EXERCISE  H.     CHEMICAL  AND  PHYSICAL  CHANGES. 
§§5  to  .12. 

7.  Upon  a  small  piece  of  quicklime  put  a  few  drops  of  water. 
Carefully  describe  all  changes  which  take  place. 

8.  Examine  a  piece  of  marble.     Notice  whether  it  is  hard  or  soft. 
From  a  piece  of  glass  tubing  of  about  |  inches  internal  diameter  cut 
off  a  piece  about  4  inches  long  by  making  a  mark  across  it  with  a 
triangular  file,  and  then  seizing  it  with  both  hands,  one  on  each  side 
of  the  mark,  pulling,  and  at  the  same  time  pressing  slightly  as  if  to 
break  it.    Clean  and  dry  it,  and  hold  one  end  in  the  flame  of  a  labora- 
tory burner  until  it  melts  together.    During  the  melting  twirl  the  tube 
constantly  between  the  finger  and  thumb  so  that  the  heat  may 
act  uniformly  upon  it. 

Heat  a  piece  of  the  marble  in  this  tube. 

Does  it  change  in  any  way  ? 

Will  it  dissolve  in  water  ? 

Treat  a  small  piece  with  dilute  hydrochloric  acid.    What  takes  place? 


LABORATORY   EXERCISES.  9 

After  the  action  has  continued  for  about  half  a  minute  insert  a 
lighted  match  into  the  upper  part  of  the  tube.  Does  the  match  con- 
tinue to  burn  ?  Does  the  substance  in  the  tube  burn  ? 

Is  the  invisible  substance  in  the  upper  part  of  the  tube  ordinary  air  ? 
How  do  you  know  ? 

Does  the  solid  substance  disappear  ? 

In  order  to  tell  whether  it  has  been  changed  chemically  the  hydro- 
chloric acid  must  be  driven  off.  This  can  be  done  by  boiling,  when  it 
passes  off  in  the  form  of  vapor,  just  as  water  does,  and  then  whatever 
is  in  solution  will  remain  behind.  For  this  purpose  put  the  solution 
in  a  small,  clean  porcelain  evaporating  dish,  and  heat  slowly  till  the 
liquid  has  disappeared. 

After  the  liquid  has  evaporated  and  the  substance  in  the  evaporat- 
ing dish  is  dry,  examine  it  and  carefully  compare  its  properties  with 
those  of  the  substance  which  was  put  into  the  test  tube.  Is  it  the 
same  substance  ?  Is  it  hard  or  soft  ?  Does  it  change  when  heated  in 
a  tube? 

Is  there  an  appearance  of  bubbling  when  hydrochloric  acid  is  poured 
on  it  ?  Does  it  dissolve  in  water  ?  Does  it  change  when  allowed  to  lie 
in  contact  with  the  air  ? 

In  order  to  learn  whether  a  substance  is  soluble  in  water  proceed  as 
follows :  Put  a  piece  about  the  size  of  a  pea  in  a  test  tube  with  pure 
water.  Thoroughly  shake,  and  .then,  as  heating  usually  aids  solution, 
boil.  Now  pour  off  a  few  drops  of  the  liquid  on  a  watch  glass,  and,  by 
gently  heating,  cause  the  water  to  pass  off  as  steam.  If  there  is  any- 
thing solid  in  solution,  it  will  be  left  on  the  platinum  foil  or  watch 


9.  Examine  crystals  of  salt.     Note  whether  they  are  hard  or  soft. 
Put  into  a  test  tube  about  5  g.  of  the  salt  and  just  cover  with  water. 
Heat  till  the  salt  dissolves.     Taste  a  drop  of  the  liquid.     Have  the  salt 
particles  been  divided  by  dissolving  ?    How  do  you  know  ? 

Arrange  a  filter  paper  (ask  the  instructor  how)  and  filter  the  solu- 
tion. Taste  a  drop  of  the  filtrate.  Has  the  salt  passed  through  the 
paper? 

Dilute  a  portion  of  the  liquid  with  an  equal  volume  of  \\Tater,  and 
shake  vigorously.  Have  the  salt  particles  been  further  divided  ? 

Evaporate  the  remainder  of  the  liquid  as  directed  in  Exp.  8.  Com- 
pare the  residue  with  the  original  salt. 

10.  Add  a  little  hydrochloric  acid,  HC1,  to  about  5  cc.  of  a  solution 
of  lead  nitrate,  Pb(N03)2.     The  hydrochloric  acid  and  lead  nitrate 


10  LABORATORY  EXERCISES. 

have  been  changed  to  lead  chloride  and  nitric  acid.     The  lead  chloride 
is  insoluble  and  is  therefore  thrown  down  (precipitated). 

Indicate  which  of  the  above  experiments  represent  physical  and 
which  chemical  changes. 

EXERCISE  III.    AIR  AND  OXYGEN. 
§§  12  to  19. 

11.  In  porcelain  crucibles  carefully  ignite  weighed  amounts  of  (a) 
zinc  dust,  stirring  occasionally;  (b)  copper  filings.1    Weigh  again. 

Was  anything  given  off  or  absorbed  ?  Could  it  have  come  from  the 
crucible  ?  From  the  gas  ?  From  whence  then  ?  Red  oxide  of  mer- 
cury was  thus  obtained  by  early  experimenters.  It  is  mercury  rust 
(just  as  the  above  are  zinc  rust  and  copper  rust,  respectively). 

12.  Heat  a  gram  of  red  oxide  of  mercury  in  an  ignition  tube  made 
as  in  Exp.  8,  holding  the  tube  nearly  horizontally.    Weigh  the  tube 
before  and  after  the  heating. 

What  evidence  is  there  of  any  change  ?  What  is  deposited  on  the 
tube  ?  Whence  did  it  come  ? 

During  the  heating  insert  into  the  tube  a  splinter  of  wood  with  a 
spark  on  the  end.  What  follows  ? 

Take  it  out  and  put  it  back  a  few  times.  Is  there  any  difference 
between  the  character  of  the  burning  in  the  tube  and  out  of  it  ? 
What  difference?  What  causes  the  difference? 

How  do  you  know  that  the  red  substance  which  you  put  into  the 
tube  has  been  changed  ?  How  could  you  collect  the  gas  given  off  ? 

Is  the  change  above  an  example  of  analysis  or  of  synthesis  ?  How 
do  you  explain  the  loss  in  weight? 

13.  How  to  collect  a  gas  over  water.    Gases  insoluble  in  water  can  be 
collected  over  that  liquid.     To  collect  gas  in  this  way,  fill  with  water 
the  vessel  to  be  filled  with  gas.     Place  over  its  mouth  a  glass  plate, 
and,  holding  the  plate  firmly  over  the  mouth  of  the  vessel  invert  it 
with  the  mouth  under  water,  in  the  pneumatic  trough. 

Does  the  vessel  remain  filled  with  water  ?    Why  ? 
Place  the  end  of  a  glass  tube  under  the  inverted  cylinder,  and  blow 
gently  through  the  tube. 
Explain  what  happens. 

i  These  should  have  been  freed  from  oil  by  means  of  ether,  and  carefully 
dried. 


LABORATORY   EXERCISES.  11 

14.  Arrange  an  apparatus  as  shown  in  the  model.    In  the  flask  (or 
in  a  large  test  tube)  put  4  to  5  g.  potassium  chlorate,  and  gently  heat 
by  means  of  a  lamp.    When  the  gas  comes  off  freely  bring  an  inverted 
cylinder  filled  with  water  over  the  end  of  the  tube,  and  let  the  bubbles 
of  gas  rise  in  the  cylinder.     Confine  the  gas  by  placing  a  glass  plate 
over  the  mouth  of  the  vessel  and  inverting  it.     Insert  into  it  a  stick 
with  a  spark  on  its  end. 

AVhat  takes  place?    Is  the  gas  contained  in  the  vessel  ordinary  air? 
What  caused  the  chemical  change  in  this  case  ?    In  what  respects  is 
this  chemical  change  like  that  in  the  last  experiment? 
Which  of  the  above  illustrates  the  combustion  of  fuel  ? 

15.  Make  a  deflagrating  spoon  by  hollowing  out  the  end  of  a  piece 
of  crayon  and  attaching  it  to  a  wire.    In  the  spoon  thus  prepared 
place  a  small  piece  of  roll  brimstone  (or  a  little  sulphur)  and  allow  it 
to  burn  in  the  air.     Notice  the  odor  of  the  fumes. 

Now  set  fire  to  another  small  portion  and  introduce  it  in  the  spoon 
into  one  of  the  vessels  containing  oxygen.  Notice  the  odor  of  the 
fumes  given  off.  Do  they  appear  to  be  the  same  as  those  given  off 
when  the  burning  takes  place  in  the  air? 

Try  a  bit  of  phosphorus  in  a  similar  manner.  Compare  the  pro- 
ducts formed  when  carbon,  sulphur,  phosphorus,  etc.,  are  burned  in 
air  and  in  oxygen. 

EXERCISE  IV.    NITROGEN. 

16.  Float  a  small  evaporating  dish  on  water  contained  in  a  pneu- 
matic trough.     Put  into  the  dish  a  small  piece  of  phosphorus,1  and  set 
fire  to  it  by  means  of  a  hot  wire.     Quickly  place  a  jar  over  it  on  a 
support  which  will  prevent  the  jar  from  sinking  more  than  an  inch  or 
two  in  water.    Why  is  the  air  at  first  forced  out  of  the  vessel  ?    Why 
does  the  air  afterward  rise  in  the  vessel  ? 

After  the  burning  has  stopped,  and  the  vessel  has  cooled  down, 
about  what  proportion  of  the  air  is  left  in  the  vessel? 

Cover  the  mouth  of  the  jar  with  a  glass  plate  and  turn  it  mouth 
upward.  Try  the  effect  of  introducing  one  after  the  other  several 
burning  bodies  into  the  gas,  as,  for  example,  a  piece  of  sulphur,  etc. 

Explain  all  that  you  have  seen. 

17.2  Put  in  a  side-neck  test  tube  about  2  g.  of  ammonium  chlo- 
ride, NH4C1.,  3  g.  of  sodium  nitrite,  NaN02,  and  moisten  with  a  few 

1  Phosphorus  should  always  be  cut  under  water,  and  he  handled  with  for- 
ceps, never  with  the  fingers. 

2  To  the  Instructor.    The  following  may  be  substituted  for  16. 


12  LABORATORY   EXERCISES. 

drops  of  water.  Apply  gentle  heat  and  collect  the  gas  over  water. 
Try  the  same  combustion  experiments  as  indicated  in  Exp.  16. 

Nitrogen  may  also  be  generated  by  passing  air  over  red  hot  copper,  Cu. 

What  compound  would  be  formed  f 

Compare  the  properties  of  oxygen  and  nitrogen. 

From  the  percentage  of  nitrogen  in  the  air  about  1  per  cent,  must 
be  subtracted  for  the  recently  discovered  element  argon,  A. 

17  (a).  Into  a  flask  fitted  with  delivery  tube  and  thistle  funnel  put 
50  cc.  of  reagent  ammonia.  Rub  about  20  g.  of  bleaching  powder 
into  a  thin  paste  with  water,  and  add  it  gradually  through  the  funnel 
while  the  flask  is  being  gently  heated.  Collect  the  gas  over  water. 

EXERCISE  V.  EFFECT  OF  TEMPERATURE  ON  VOLUME  OF  GASES. 

18.  Fit  a  small  flask  with  a  glass  tube  about  20  cm.  long  by  means 
of  a  perforated  rubber  stopper.     Support  the  tube  or  flask  in  a  vertical 
position,  flask  uppermost,  in  a  vessel  of  colored  water.     Apply  heat  to 
the  flask  by  means  of  the  burner. 

Note  the  result.    What  escapes  through  the  tube?    Why? 

CAUTION.  —  Do  not  continue  the  heating  too  long  or  the  flask  may  be  broken 
by  the  inrush  of  water  on  the  removal  of  the  heat. 

Why  does  the  water  rush  in  after  the  removal  of  the  heat  ?  State 
the  effect  of  heat  on  the  volume  of  a  gas. 

Suppose  the  flask  to  have  been  tightly  stoppered,  and  heat  to  have 
been  applied  in  a  like  manner,  how  would  the  pressure  under  which 
the  gas  existed  have  been  affected? 

COEFFICIENT  OF  EXPANSION  OF  GASES. 

19.  For  the  Instructor.    Into  a  glass  tube  about  20  cm.  long  and  of  1 
mm.  bore,  introduce  for  an  index  about  5  mm.  of  mercury  by  placing 
one  end  of  the  tube  in  a  bottle  of  mercury  and  holding  the  finger 
on  the  stopper  end  of  the  tube  as  it  is  removed  from  the  bottle.     In- 
cline the  tube  so  as  to  work  the  mercury  column  to  about  the  center 
of  the  tube.     Now,  holding  the  tube  in  a  horizontal  position,  close  one 
end  of  it  by  holding  in  a  Bunsen  flame.     Fasten  a  chemical  thermom- 
eter to  the  tube  by  means  of  wire  or  string  so  it  may  be  used  as  one 
of  the  divisions  for  measuring  the  length  of  the  column  of  confined 
air.     Have  ready  a  long  shallow  dish  containing  ice  water.     Place  the 
apparatus  in  the  dish  of  ice  water  in  as  near  a  horizontal  position  as 


LABORATORY   EXERCISES.  13 

possible,  and  determine  the  length  of  the  inclosed  air  column  for  10°  C., 
20°  C.,  and  30°  C.,  heating  the  tube  by  the  gradual  addition  of  hot 
water  to  the  water  in  the  dish. 

From  the  data  obtained  compute  the  expansion  per  degree  of  the 
air  per  unit  of  volume  as  measured  at  0°  C.,  between  0°  C.  and  10°  C., 
10°  C.  and  20°  C.,  etc.  If  the  work  is  carefully  done  the  result  should 
be  close  to  .00366,  or  5}y  of  the  bulk  per  1°  C. 

At  what  temperature  would  1  cc.  of  gas  become  2  cc.,  the  pressure 
being  constant  ?  What  would  be  the  volume  of  1  cc.  measured  at  0°  C., 
when  cooled  to  —  273°  C.?  What  temperature  is  called  absolute  zero? 

NOTE.  —  The  student  must  understand  that  these  quantitative  experiments 
require  the  utmost  care,  and  even  then,  with  the  crude  apparatus  used,  and 
the  limited  experience  in  doing  such  work,  only  approximate  results  can  be 
expected.  The  above  experiments  are  intended  to  illustrate  the  fact  that 
gases  expand  with  the  increase  of  heat,  and  vice  versa,  and  therefore  vary  in 
the  pressure  they  exert;  and  also  the  LAW  OF  CHARLES:  "All  true  gases, 
when  heated,  expand  5^3  of  their  volume,  measured  at  OP  C.,for  each  increase 
ofl°C." 

Example  i.  Twenty  cubic  centime'ters  of  hydrogen  were  measured 
at  15°  C.  and  heated  to  35°  C.  What  was  the  new  volume  ? 

Example  2.  Five  hundred  cubic  centimeters  of  nitrogen  measured 
at  27°  C.  would  become  how  many  when  cooled  to  —10°  C.? 

EXERCISE  VI.  EFFECT  OF  PRESSURE  ON  VOLUME  OF  GASES. 

20.  For  the  Instructor.  Procure  a  tube  shaped  like  the  letter  J,  the 
short  arm  being  closed  and  about  25  cm.  long,  and  the  long  arm  at 
least  85  cm.  in  length  and  open  at  the  top.  There  will  be  less  liability 
to  accident  and  consequent  loss  of  mercury  if  the  tube  be  fastened  in 
an  upright  position,  either  by  attaching  it  to  a  board  set  in  a  block 
for  a  base,  or  by  tying  it  securely  to  the  standard  of  a  ring  stand. 
Divide  the  short  arm  into  four  equal  parts,  indicating  the  divisions 
by  means  of  rubber  bands  or  gummed  paper ;  also  make  a  mark  near 
the  base  of  the  long  arm  exactly  opposite  the  lowest  mark  on  the  short 
arm,  which  must  be  above  the  bend  of  the  tube.  Now  pour  mercury 
into  the  tube  through  a  small  funnel  till  the  surfaces  in  the  two  arms 
stand  at  the  lowest  division  in  each  arm.  We  have  now  entrapped 
some  air,  and  the  tension  (pressure)  in  the  short  arm  is  the  same  as 
in  the  long  arm.  How  is  this  indicated  ?  Under  how  much  pressure 
per  square  inch  is  it  ?  Now  pour  mercury  into  the  long  arm,  carefully 
inclining  the  tube  so  as  to  include  as  little  air  as  possible,  till  the  space 


14  LABORATORY   EXERCISES. 

in  the  short  arm  has  been  reduced  one  half.  Measure  the  height  of 
the  column  of  mercury  in  the  long  arm  above  the  mark  near  the  base, 
and  compare  it  with  the  reading  of  the  barometer  at  the  time  of  the 
experiment.  Remembering  what  was  the  pressure  of  the  confined  air 
in  the  beginning,  what  is  its  tension  now  ?  Under  how  many  atmos- 
pheres' pressure  is  it?  How  is  this  shown?  How  does  its  former 
volume  compare  with  that  occupied  now?  Doubling  the  pressure  has 
how  affected  the  volume  occupied  ?  How  has  it  affected  the  tension 
of  the  gas? 

If  the  length  of  the  tube  will  allow,  add  as  much  more  mercury  to 
the  long  arm.  Under  how  much  pressure  is  the  confined  air  now? 
How  does  its  volume  compare  with  the  original  volume?  Under 
about  how  much  pressure  per  square  inch  is  it  ? 

EFFECT  OF  DECREASING  THE  PRESSURE. 

21.  In  the  previous  experiment  the  pressure  on  the  confined  air  was 
increased.  Let  us  now  reverse  the  condition  and  ascertain  the  effect 
of  diminishing  the  pressure. 

Fill  a  barometer  tube  to  within  about  10  cm.  of  the  top.  Place  the 
forefinger  over  the  open  end,  thus  inclosing  a  certain  volume  of  air. 
Under  what  tension  is  the  inclosed  air  ? 

Now  invert  the  tube  and  allow  the  air  to  rise  to  the  top.  Keeping 
the  finger  over  the  end  of  the  tube,  measure  the  length  of  the  air 
column.  Place  the  mouth  of  the  tube  under  mercury  contained  in 
a  porcelain  mortar  and  remove  the  finger,  and  the  mercury  column 
will  fall,  for  the  atmosphere  outside  cannot  support  the  atmospheric 
pressure  of  the  column  plus  the  weight  of  the  mercury  column.  As 
the  mercury  falls  the  imprisoned  air  expands  and  presses  less  on  the 
mercury  column,  and  a  point  is  soon  reached  at  which  the  pressure 
outside  and  inside  the  tube  is  equal.  The  inclosed  air  is  under  the 
atmospheric  pressure  (how  many  cm.  of  mercury  ?)  less  the  pressure 
due  to  the  column  of  mercury  in  the  tube. 

Call  the  original  volume  of  inclosed  air  Vl  and  the  final  volume  V2 ; 
the  original  tension  Pl  and  the  final  tension  P2;  the  results  should 
give  very  close  to  the  following  equation,  provided  the  work  has  been 
carefully  done. 

Vj  x  P!  =  V2  x  P2. 

Repeat  the  experiment  with  a  new  volume  of  inclosed  air,  and  see  if 
the  results  agree. 


LABORATORY  EXERCISES.  15 

Express  as  a  law  the  relation  between  volume  and  pressure  as 
exhibited  by  these  experiments.  This  is  known  as  BOYLE'S  LAW. 
The  law  is  not  rigorously  correct  for  all  gases,  the  variations  depend- 
ing on  the  kind  of  gas  used.  It  is  sufficiently  exact  for  all  practical 
purposes,  however.  The  law  does  not  hold  for  gases  when  the 
measurement  is  taken  too  near  their  point  of  liquefaction. 

Example  i.  Twenty  cubic  centimeters  of  hydrogen  were  measured 
at  760  mm.  pressure  and  afterward  the  pressure  changed  to  768  mm. 
What  volume  did  the  gas  then  occupy  ? 

Example  2.  Five  hundred  cubic  centimeters  of  nitrogen  measured 
at  768  mm.  would  become  how  many  cubic  centimeters  at  755  mm.  ? 

EXERCISE  VII.    WEIGHT  AND  DENSITY  OP  AIR. 

22.  Fit  a  stout  Florence  flask  of  about  a  liter  capacity  with  a  one- 
hole  rubber  stopper  through  which  passes  a  tightly  fitting  glass  tube 
carrying  a  stiff  pinchcock.  Make  the  apparatus  air-tight  by  smear- 
ing the  points  with  vaseline,  if  necessary.  Be  sure  the  flask  is  clean 
and  dry. 

First,  weigh  the  apparatus  full  of  air  and  record  the  weight. 

Second,  open  the  pinchcock  and  by  means  of  an  air  pump  exhaust 
the  air  from  the  flask  as  thoroughly  as  possible.  Close  the  pinch- 
cock, reweigh  the  apparatus,  and  note  the  loss  in  weight. 

Third,  measure  the  volume  of  air  that  was  removed,  as  follows : 
Place  the  rubber  tube  bearing  the  pinchcock  well  beneath  the  surface 
of  water  in  a  pneumatic  trough.  Open  the  pinchcock,  keeping  the  tube 
well  under  water  all  the  time.  The  water  will  rush  into  the  flask. 
Why  ?  Make  the  level  of  the  water  inside  and  outside  the  flask  equal 
by  lowering  or  raising  the  flask  in  the  water.  Close  the  pinchcock, 
remove  the  apparatus  from  the  water,  and  make  it  dry.  Ascer- 
tain the  number  of  cubic  centimeters  of  water  in  the  flask  by  weigh- 
ing on  a  balance,  remembering  that  1  g.  equals  1  cc.  of  water. 
From  the  data  thus  secured  compute  the  weight  of  one  cubic  centi- 
meter of  air  under  the  conditions  of  the  experiment,  or  the  density 
of  air. 

Define  density. 

REMARK.  — The  term  density  is  often  used  as  synonymous  with  specific 
gravity,  but  the  latter  expresses  more  properly  the  number  of  times  a  given 
volume  of  a  certain  substance  is  heavier  than  some  other  substance  taken  as  a 
standard,  while  density  refers  to  the  weight  of  a  unit  volume. 


16  LABORATORY   EXERCISES. 

Referring  to  19,  20,  and  21,  how  would  the  density  of  air  have  been  affected 
had  you  made  the  determination  at  0°  C,  and  760  mm.  pressure,  which  are  the 
standard  conditions  ? 

EXERCISE  VIII.    WATER. 

§§  34  to  37. 

23.  Hold  a  dry  cold  beaker  for  an  instant  over  a  Bunsen  burner 
flame.     Is  any  water  deposited  on  it  ? 

In  a  dry  test  tube  heat  gently  a  small  piece  of  wood.     What  evi- 
dence do  you  obtain  that  water  is  given  off? 
Try  the  same  with  a  little  sugar  in  a  test  tube. 

24.  For  the  Instructor.1     Roll  a  piece  of  metallic  sodium  in  a  piece 
of  wire  gauze  and  by  means  of  a  pair  of  tongs  thrust  the  cage  under 
the  mouth  of  an  inverted  receiver  containing  water.     After  the  action 
has  ceased  close  the  mouth  of  the  receiver,  and  turn  it  uppermost. 
Remove  the  plate  and  apply  a  lighted  taper  to  the  gas. 

Does  it  act  like  oxygen  ?    How  does  it  differ  ? 

25.  For  the  Instructor.1    In  a  retort  of  convenient  size  distil  water 
before  the  class.     Call  attention  to  the  colorless  nature  of  the  steam. 
Place  a  few  drops  of  the  distilled  water  on  a  platinum  foil  and  evapor- 
ate.    Then  do  the  same  with  a  few  drops  of  water  remaining  in  the 
retort. 

*What  effect  has  distillation  on  the  purity  of  water  ? 
Does  well  water  contain  solids  in  solution  ? 
Why  has  distilled  water  not  an  agreeable  taste? 

26.  Dissolve  2  or  3  g.  of  common  salt  in  distilled  water.     Evapor- 
ate the  solution  slowly  to  dryness  and  compare  the  substance  ob- 
tained with  the  original  salt  as  to  appearance,  taste,  and  crystalline 
form. 

Dissolve  5  g.  of  sodium  carbonate,  Na2C03,  in  dilute  hydrochloric 
acid,  HC1.  Evaporate  the  solution  as  before,  and  compare  the  resi- 
due with  the  original  salt  as  to  appearance  and  taste,  and  by  treating 
with  hydrochloric  acid. 

Explain  the  difference  between  physical  and  chemical  solution. 

27.  Dissolve  some  ordinary  alum  in  water  (f  ounce  alum  to  100 
cc.  water)  by  the  aid  of  heat.     Filter  through  a  plaited  filter  and 
allow  the  filtered  solution  to  cool. 

What  takes  place? 

1  These  may  be  performed  in  the  lecture  room. 


LABORATORY    EXERCISES.  17 

Which  will  hold  in  solution  the  greater  amount  of  substance  —  hot 
water  or  cold  water  ? 

Define  crystallization. 

28.  Pour  off  the  above  liquid  and  place  a  few  of  the  crystals  on  a 
piece  of  dry  filter  paper.  After  the  water  is  all  absorbed  from  them 
and  they  appear  dry,  put  them  in  a  dry  test  tube  and  heat  gently. 

What  evidence  have  you  that  water  is  contained  in  the  crystals  ? 

29.1  Some  bodies  give  up  their  water  of  crystallization  simply  on 
contact  with  the  air.  Such  bodies  are  said  to  be  efflorescent. 

Put  some  crystals  of  sodium  sulphate,  Na2S04  •  10  H20,  on  a  watch 
crystal  and  expose  them  for  an  hour  or  more  to  the  air  of  the  room. 
Try  also  calcium  chloride.  Substances  acting  like  the  latter  are  said 
to  deliquesce. 

Define  efflorescence ;  also  deliquescence. 

30.  Take  a  glass  tube  3  or  4  cm.  in  diameter,  and  close  one  end 
with  a  plug  of  plaster  of  Paris  1  or  2  cm.  thick.    Set  the  tube  aside 
to  dry  until  the  next  exercise. 

EXERCISE  IX.  ANALYSIS  OF  WATER. 

31.  For  the  Instructor.     Fill  a  water  decomposition  apparatus  with 
water  acidulated  with  one-tenth  its  volume  of  sulphuric  acid.    Connect 
the  platinum  electrodes  of  the  apparatus  with  the  wire  from  a  Bunsen 
battery.     Start  the  current,  and  bubbles  will  collect  in  each  of  the  two 
arms.     Compare  the  volume  of  the  two  gases  which  collect  in  the  two 
tubes  in  a  given  time.     When  one  of  the  tubes  is  full  hold  a  lighted 
match  above  the  tube,  open  the  stopcock,  and  ignite  the  gas.     Is  the 
gas  air?    How  do  you  know?    Does  the  gas  burn  ? 

Immerse  a  glowing  splinter  in  the  gas  in  the  other  tube.  Is  the 
action  the  same  as  in  the  former  case?  Is  the  gas  in  this  tube  air? 
Is  it  steam  ?  Does  it  fume  ?  How  does  it  act  toward  the  splinter? 

In  this  experiment  the  water  has  been  separated  into  its  constituent 
parts  —  analyzed. 

Define  analysis. 

The  reverse  of  this  would  be  synthesis. 

LAW  OF  DEFINITE  PROPORTIONS. 

32.  For  the  Instructor.     Fill  a  eudiometer  tube  with  mercury  and 
invert  it  over  a  dish  of  mercury.     Now  from  a  gasometer  admit  to  the 
tube,  which  should  be  held  in  a  slanting  position,  10  cc.  of  hydrogen 

1  This  experiment  should  be  arranged  at  the  beginning  of  the  exercise. 

LAB.    EX.  — 2 


18  LABOEATOEY  EXBECISES. 

which  is  first  passed  through  a  wash  bottle  containing  a  solution  of 
2  g.  of  caustic  potash  in  10  cc.  of  water,  and  then  through  a  second 
wash  bottle  containing  strong  sulphuric  acid.  Bring  the  tube  to  an 
upright  position  and  secure  the  following  data:  (a)  volume  of  gas, 
(6)  temperature,  (c)  reading  of  barometer,  (d)  length  in  millimeters 
of  the  mercury  column  in  the  tube. 

Again  slant  the  tube,  and  admit  in  a  similar  manner  about  8  cc. 
of  oxygen  from  the  gasometer.  Now  bring  the  tube  to  an  upright 
position,  and  again  make  readings  as  above. 

Press  the  eudiometer  tightly  against  a  leather  or  rubber  washer  on 
the  bottom  of  the  trough  and  clamp  tightly  in  position.  From  an 
induction  coil  pass  a  spark  through  the  gases  by  connecting  wires 
from  the  coil  to  the  electrodes  of  the  tube. 

After  the  tube  has  cooled  raise  slightly  and  make  the  same  readings 
as  in  the  two  instances  above. 

Calculate  all  the  volumes  to  the  standard  conditions  as  given 
under  Remarks,  Exp.  22. 

What  became  of  the  extra  volume  of  oxygen?  Test  the  gas 
remaining  in  the  tube  for  oxygen,  as  under  Exp.  12. 

Repeat  the  experiment,  using  an  excess  of  hydrogen.1 

In  what  proportion  did  the  two  gases  unite  each  time  ? 

What  lesson  may  be  learned  from  these  two  experiments  as  to  the 
proportion  in  which  elements  unite  volurnetrically  ? 

EXERCISE  X. 

33.  Into  a  small  tared  beaker  weigh  just  5  g.  of  sodium  carbonate 
crystals  (they  must  not  have  effloresced)  and  dissolve  them  in  water. 
Now  add  hydrochloric  acid  little  by  little  as  long  as  any  effervescence 
takes  place.  Now  evaporate  the  water  and  ascertain  the  weight  of  the 
salt  remaining. 

Care  must  be  exercised  toward  the  last  of  the  drying  lest  loss  occur 
from  spattering. 

Now  repeat  in  exactly  the  same  manner,  except  that  after  all 
effervescence  has  ceased  an  excess  of  the  acid  is  added.  Compare  the 
weight  of  salt  obtained  in  this  case  with  the  first.  Did  any  different 
amount  of  the  acid  unite  in  the  second  case  from  that  in  the  first  ? 
What  is  effervescence  ? 
.  State  your  observation  in  the  form  of  a  law. 

1  These  gases  should  be  admitted  through  a  tube  drawn  out  to  a  small  jet. 


LABORATORY   EXERCISES.  19 

34.  Conservation  of  Matter.  One  of  the  laws  which  the  studeut 
must  have  impressed  upon  his  mind  early  in  his  chemical  study  is 
that  in  the  various  changes  which  substances  may  undergo,  nothing  is 
lost.  No  matter  is  destroyed,  nor  is  any  made.  This  law  is  the  basis 
of  all  physical  science. 

This  fact  may  be  illustrated  as  follows :  — 

Make  a  saturated  solution  of  calcium  chloride  by  adding  the  salt  to 
about  20  cc.  of  water  as  long  as  it  will  dissolve.  In  a  like  manner 
prepare  an  equal  amount  of  a  saturated  solution  of  sodium  sulphate. 
Fill  a  test  tube  a  little  less  than  half  full  of  the  first  solution,  and  in 
another  tube  place  an  equal  amount  of  the  second  solution.  Place  the 
two  tubes  in  a  beaker  to  keep  them  from  overturning,  and  determine 
the  combined  weight.  Pour  one  solution  into  the  other,  and,  after 
shaking  and  observing  the  change,  again  ascertain  the  weight.  Has 
weight  either  been  gained  or  lost  in  this  operation?  State  in  the 
form  of  a  law  the  fact  here  illustrated. 


EXERCISE  XI.    MULTIPLE  PROPORTIONS. 

35.  For  the  Instructor.  In  the  blue  flame  heat  a  small  clean  porcelain 
crucible  supported  on  a  pipestone  triangle.  Allow  it  to  cool  in  a  desic- 
cator containing  either  sulphuric  acid,  or  some  pieces  of  calcium  chloride. 
When  perfectly  cold  weigh  the  crucible  on  a  delicate  balance,  and  record 
the  weighing.  Place  in  the  crucible  a  layer  of  dry  copper  oxide. 
Again  obtain  the  weight.  After  placing  the  crucible  again  on  the 
pipestone  support,  cover  it  with  a  cover  having  in  the  middle  a  small 
hole.  Connect  a  bent  porcelain  tube  with  a  Kipp's  hydrogen  gen- 
erator (or  a  gasometer),  and  turn  on  a  current  of  hydrogen,  directing 
it  into  the  crucible  by  means  of  the  bent  porcelain  tube  through 
the  hole  in  the  cover.  Heat  the  crucible  with  this  current  of  hydro- 
gen passing  in  for  about  ten  minutes.  Allow  the  crucible  to  cool  in 
the  atmosphere  of  hydrogen,  after  which  weigh.  Take  the  difference 
in  the  two  weighings  as  the  weight  of  oxygen  given  up  by  the  given 
weight  of  the  copper  oxide  used.  Calculate  the  weight  of  copper 
combined  with  eight  parts  of  oxygen. 

Repeat  the  experiment,  using  red  cuprous  oxide  in  place  of  the 
black  oxide,  and  calculate  in  the  same  manner.  Compare  the  two 
proportions  with  each  other.  What  does  the  experiment  show? 
State  the  law  of  multiple  proportions.  Find  in  the  text-book  other 
examples  of  multiple  proportions. 


20  LABORATORY   EXERCISES. 

EXERCISE  XII.    HYDROGEN. 
§§  37  to  41. 

The  instructor  will  have  prepared  a  gasometer  filled  with  the  gets  from 
which  it  can  be  obtained  by  the  students  for  studying  its  properties  in 
Exp.  38. 

36.  Fill  the  plugged  tube  prepared  in  Exp.  30  with  hydrogen 
and  set  it  upright  in  a  glass  of  water.     Examine  it  from  time  to 
time  during  the  laboratory  period.     Describe  the  result. 

37.  Into  a  cylinder  or  test  tube  put  a  few  pieces  of  granulated  zinc, 
and  pour  upon  it  enough  ordinary  hydrochloric  acid  to  cover  it. 

After  the  action  has 'continued  for  a  minute  or  two,  apply  a  lighted 
match  to  the  mouth  of  the  vessel. 

Describe  in  full. 

Perform  the  same  experiments,  using  sulphuric  acid  which  has  been 
diluted  with  four  times  its  volume  of  water.1 

What  is  the  result? 

Try  iron  and  sulphuric  acid. 

From  what  does  the  hydrogen  come  ? 

Ascertain  if  same  result  is  obtained  by  substituting  zinc  oxide  for 
zinc. 

38.  Carefully  lift  from  the  water  pan  a  bottle  completely  full  of 
hydrogen. 

Slowly  carry  the  bottle,  the  mouth  held  downward,  to  a  burning 
splinter  of  wood,  and  depress  the  bottle  over  this  flame.  After 
observing  what  happens,  withdraw  the  taper  slowly. 

Does  the  hydrogen  burn  at  the  surface  or  at  the  end  of  the  splinter  ? 

Does  hydrogen  support  combustion  ? 

What  gathers  on  the  inside  of  the  jar  when  the  hydrogen  burns  ? 

1  To  dilute  ordinary  concentrated  sulphuric  acid  with  water,  the  acid  should 
be  poured  slowly  into  the  water  while  the  mixture  is  constantly  stirred.  If  the 
water  is  poured  into  the  acid,  the  heat  evolved  at  the  places  where  the  two 
liquids  come  in  contact  with  each  other  may  be  so  great  as  to  convert  the 
water  into  steam  and  cause  the  strong  acid  to  spatter. 

In  experimenting  with  hydrogen,  no  light  should  ever  be  brought  into  con- 
tact with  the  contents  of  the  bottle  in  which  it  is  generated,  or  with  any  large 
quantity  of  the  gas,  until  the  purity  of  the  sample,  or  rather  its  nonexplosive 
character,  has  been  demonstrated  by  applying  to  a  very  small  volume  of  the 
gas  the  test  above  described, 


LABORATORY   EXERCISES.  21 

EXERCISE  XIII.    COMPOUNDS  OF  OXYGEN  AND  NITROGEN. 

§§  57  to  64. 

39.  Place   a  small   quantity  of  ammonium  nitrate,  NH4N03,  in  a 
test  tube,  and  heat.     Hold  a  piece  of  cool  glass  near  the  mouth  of  the 
tube,  and  note  what  collects  upon  it. 

What  is  the  first  change  that  takes  place  ?     What  is  the  next  ? 
Complete  the  equation,  NH4N03  =  2H20  +? 

40.  Collect  over  water  a  jar  of  nitrogen  monoxid,  N20,  made  by 
heating  about  5  g.  of  ammonium  nitrate  in  a  side-neck  tube.      Do 
not  heat  higher  than  is  necessary  to  secure  a  regular  evolution  of 
the  gas. 

Observe  its  taste,  odor,  and  effect  upon  a  burning  stick. 
Explain  the  action  of  the  burning  gas  toward  the  burning  stick. 

41.  NITRIC  OXIDE.     For  the  Instructor.     Fit  a  flask  with  both  de- 
livery and  safety  tubes.     Into  this  flask  put  pieces  of  copper  turnings. 
Cover  them  with  water.     Now  slowly  add  ordinary  concentrated  nitric 
acid.    When  enough  acid  has  been  added  gas  will  be  given  off.     If 
the  acid  is  added  quickly  the  evolution  of  gas  takes  place  too  rapidly, 
so  that  the  liquid  is  forced  out  of  the  flask  through  the  funnel  tube. 
This  can  be  avoided  by  not  being  in  a  hurry. 

Do  not  inhale  the  gas.  Perform  the  experiments  with  nitric  oxide 
where  there  is  a  good  draught. 

What  is  the  color  of  the  gas  in  the  flask  at  first  ? 

What  is  it  after  the  action  has  continued  for  a  short  time  ? 

Collect  over  water  two  or  three  vessels  full. 

Balance  the  following  equation  —  3Cu  +  8HN03=  3Cu(N03)2  + 
H20+N0. 

42.  Turn  one  of  the  cylinders   of  nitric  oxide  with   the   mouth 
upward  and  uncover  it.     What  takes  place  ? 

What  element  in  the  air  is  most  likely  to  be  the  cause  of  this 
change  ?  Does  the  gas  after  exposure  to  the  air  resemble  that  in  the 
generating  flask  at  first  ? 

Explain  the  presence  of  the  colored  gas  at  the  beginning  of  Exp. 
41  and  the  fact  that  it  finally  disappeared. 

How  many  oxides  of  nitrogen  are  described  in  your  text-book  ? 

Make  a  table  of  them,  showing  the  relation  N  to  0,  by  weight. 

What  fundamental  law  of  chemical  action  may  be  derived  from  a 
consideration  of  these  compounds  ? 


22  LABORATORY  EXERCISES. 

EXERCISE  XIV.    NITRIC  ACID. 

§§  64  to  70. 

43.  Instructor  prepare  nitric  acid,  HN03,  as  per  S.  &  L.,  p.  55, 
Exp.  33.     The  retort  used  must  be  a  glass-stoppered  one. 

Using  the  acid  thus  prepared  (if  sufficient  for  the  entire  class  has 
not  been  prepared,  then  a  portion  of  the  class  may  demonstrate  to  the 
others),  let  the  students  ascertain  the  properties  by  the  following 
experiments :  — 

44.  To  1  cc.  of  nitric  acid  of  Exp.  43  add  10  cc.  of  water.     Touch  a 
drop  of  the  mixture  to  the  tongue.     Dip  a  piece  of  blue  litmus  paper 
into  the  liquid. 

45.  With  a  stirring  rod  place  a  drop  of  the  strong  acid  on  the 
finger  nail,  and  after  a  moment  wash  it  off.     Put  a  few  pieces  of 
white  wool  or  worsted  into  a  few  drops  of  the  strong  acid  in  an 
evaporating  dish  and  warm  gently. 

What  is  the  effect  of  the  acid  on  organic  substances  ? 

46.  Try  the  effect  of  the  acid  on  a  bit  of  copper.     Place  in  each 
of  two  test  tubes  a  piece  of  zinc,  Zn.     To  one  add  dilute  hydrochloric 
acid  and  to  the  other  a  few  drops  of  strong  nitric  acid  diluted  with 
an  equal  volume  of  water.     What  difference  is  noticed  in  the  action  ? 

Try  the  same  with  iron  filings. 

How  does  nitric  acid  act  upon  metals  ? 

47.  When  nitric  acid  acts  on  metals  nitrates  are  formed :   e.g.  nitric 
acid  acting  on  copper.    In  such  cases  the  hydrogen,  H,  of  the  acid  is 
replaced  by  the  metal.    What  becomes  of  the  hydrogen  ? 

These  nitrates  are  good  oxidizing  agents.  Melt  in  a  tube  a  few 
pieces  of  potassium  nitrate,  and  then  drop  in  a  bit  of  charcoal  and 
heat. 

Repeat,  using  a  small  piece  of  roll  sulphur  instead  of  the  charcoal. 


EXERCISE  XV.    BASES,  SALTS. 
§§  70  to  72. 

48.  Test  the  acids  that  you  find  in  the  laboratory  as  to  their  effect 
upon  litmus  paper,  and  also  as  to  their  taste  after  diluting  them  with 
ten  or  fifteen  times  their  volume  of  water.  How  do  they  compare  in 
their  volume  action  ? 

What  does  this  experiment  teach  concerning  acids? 


LABORATORY   EXERCISES.  23 

49.  Into  about   20  cc.  of  water  put  a  few  drops  of  potassium 
hydrate  solution,  KOH.    Rub  a  little  of  the  KOH  solution  between 
the  fingers.     Cautiously  taste  of  the  diluted  solution.     Immerse  a 
piece  of  red  litmus  paper  in  it. 

Try  the  same  with  sodium  hydrate,  NaOH,  and  ammonium  hydrate, 
NH4OH.  What  does  this  experiment  teach  concerning  bases  ?  How 
do  the  bases  act  toward  litmus  as  compared  with  the  acids  ? 

50.  To  a  solution  of  caustic  soda,  NaOH,  add  dilute  hydrochloric 
acid  slowly,  examining  the  solution  from  time  to  time  by  means  of  a 
piece  of  paper  colored  blue  with  litmus.    As  long  as  the  solution  is 
alkaline  it  will  cause  no  change  in  the  color  of  the  paper.     The 
instant  it  passes  the  point  of  neutralization  it  changes  the  color  of 
the  paper  red.     When  this  point  is  reached,  evaporate  the  water  on 
a  water  bath  to  complete  dryness,  and  see  what  is  left.     Taste  the 
substance.     Has  it  an  acid  taste? 

Does  it  suggest  any  familiar  substance  ? 

Is  it  an  acid,  an  alkali,  or  is  it  neutral  ? 

Treat  a  little  of  the  material  with  sulphuric  acid,  H2S04,  and  note 
its  action.  Treat  a  little  common  salt,  NaCl,  in  a  similar  manner. 
How  do  the  two  substances  compare  in  action  ? 

Write  the  equation  showing  the  formation  of  this  salt. 

51.  Test  with  litmus  paper  the  reaction  of  other  salts  in  solution  ; 
for  example,  ammonium  sulphate,  (NH4)2S04,  and  potassium  chloride, 
KC1. 

What  is  the  usual  action  of  salts  toward  litmus  ? 
Try,  however,  solutions  of  sodium  carbonate,  Na2C03,  copper  sul- 
phate, CuS04,  sodium  bicarbonate,  NaHC03,  sodium  bisulphate,  NaHS04. 

EXERCISE  XVI.     AMMONIA. 

§§  72  to  78. 

52.  Take  in  one  hand  a  little  dry  quicklime,  CaO,  and  in  the  other 
an  equal  bulk  of  pulverized  ammonium  chloride,  NH4C1.    Note  that 
neither  substance  has  an  odor.    Now  rub  the  two  together  between 
the  hands,  and  carefully  note  the  odor  of  the  gas  given  off. 

53.  To  a  solution  of  ammonium  chloride  in  a  test  tube  add  a  few 
drops  of  potassium  hydroxide,  KOH.     Warm  gently,  and  note  the 
odor  of  the  fumes,  as  well  as  their  action  toward  litmus. 

Also  do  the  same  with  a  solution  of  ammonium  nitrate,  NH4N03. 
What  result  ?    Write  equations  for  each  reaction. 


24  LABORATORY   EXERCISES. 

Moisten  a  stirring  rod  with  hydrochloric  acid,  and  hold  it  in  the 
escaping  ammonia  fumes.  Try  the  same  with  nitric  acid.  The  two 
gases  unite  in  each  case  to  form  a  solid.  Write  equations  to  show  the 
union. 

What  two  classes  of  compounds  are  used  to  make  ammonia  ? 

54.  Fit  a  right-angled  bend  to  a  flask  by  means  of  a  perforated 
stopper.     Connect  to  this  bend  a  second  tube  of  the  same  shape,  direct- 
ing the  free  end  downward.     To  the  lower  end  of  this  tube  adjust  a 
glass  funnel,  inverted.     Allow  the  mouth  of  the  funnel  to  dip  under 
water  held  in  a  beaker.     Into  the  flask  put  two  parts  of  ammonium 
chloride  to  one  of  quicklime.     Moisten  slightly  and  apply  gentle  heat. 
After  the  gas  has  escaped  in  the  water  for  a  few  minutes  disconnect 
the  funnel  and  direct  the  gas  upward,  by  turning  the  tube,  into  a 
cylinder  held  over  the  end  of  the  delivery  tube. 

Into  the  gas  thus  collected  insert  a  burning  stick.  Does  the  gas 
burn  ?  Does  it  support  combustion  ?  Is  the  gas  soluble  in  water  ? 

Test  with  red  litmus  paper  the  water  into  which  the  gas  has  been 
passed.  Is  it  acid  or  is  it  alkaline ?  What  is  the  liquid? 

What  is  the  relation  between  ammonia  and  ammonium  hydroxide? 

Make  in  your  notebook  a  tabular  statement  of  the  physical  and 
chemical  properties  of  ammonia. 

55.  Place  in  a  beaker  a  little  dilute  nitric  acid ;  now  carefully  add  a 
solution  of  ammonia  (ammonium  hydroxide,  NH4OH)  until  the  acid 
is  neutralized. 

Record  how  you  tested. 

Slowly  evaporate  this  solution  to  dryness  and  examine  the  salt. 

What  is  its  name  ?    Write  equation  for  its  formation. 

What  salt  would  have  been  formed  if  sulphuric  acid  had  been  used  ? 

Write  an  equation  to  show  its  formation. 

EXERCISE  XVII.     HYDROCHLORIC  ACID,  HC1. 

§§  78  to  87. 

56.  Put  a  little  common  salt,  NaCl,  into  a  dry  test  tube,  and  pour 
on  it  a  few  drops  of  strong  sulphuric  acid,  H2S04.  Note  the  character- 
istics of  the  gas  given  off. 

Test  with  a  piece  of  blue  litmus  paper. 

Try  the  same  with  other  chlorides,  as  KC1,  NH4C1,  CaCl2. 

Do  you  obtain  HC1  in  each  case? 

Write  an  equation  for  each  reaction. 


LABORATORY   EXERCISES.  25 

57.  In  a  flask  fitted  with  a  delivery  and  a  thistle  tube  generate  HC1 
from  20  gr.  sodium  chloride,  NaCl,    15  cc.   water  and   10  cc.   sul- 
phuric acid,  H2S04.      The  contents  of  the  flask  must  be  very  gradually 
and  moderately  heated,  else  a  violent  frothing  is  liable  to  occur  which 
would  spoil  the  experiment.     In  your  notes,  describe  the  apparatus; 
and  show  the  use  of  safety  tubes.     Let  the  delivery  tube  pass  into  a 
cylinder,  keeping  it  covered  as  well  as  possible.     After  filling  three 
cylinders  by  downward  displacement  let  the  gas  pass  into  a  beaker  of 
water. 

58.  Test  a  cylinder  of  the  gas  with  a  lighted  taper;   notice  also 
the  white  fumes  which  are  formed  when  the  gas  comes  in  contact  with 
moist  air. 

59.  Moisten  a  piece  of  paper  with  ammonium  hydroxide,  NH4OH, 
and  thrust  it  into  a  cylinder  of  the  gas. 

Explain  what  takes  place,  writing  an  equation  for  the  reaction. 

60.  Invert  a  cylinder  of  the  gas,  well  covered  with  a  glass  plate,  in 
a  dish  of  water,  and  then  remove  the  cover.     Explain. 

61.  Now  test  the  liquid  in  the  beaker.     Has  it  acquired  acid  prop- 
erties ? 

Compare  the  action  of  this  liquid  with  that  marked  HC1  in  the 
reagent  bottle,  as  to  the  effect  on  litmus;  on  a  piece  of  marble;  on 
zinc  or  iron,  and  on  a  solution  of  silver  nitrate,  AgN03;  of  lead  nitrate, 
Pb(N03)2;  mercurous  nitrate,  HgN03. 

If  you  should  evaporate  to  dryness  the  solution  in  the  flask  used  in 
Exp.  57  what  would  the  residue  be  ? 

If  you  should  recover  all  the  compound  formed,  how  much  would 
you  obtain  ? 

EXERCISE  XVIII.    CHLORINE. 

62.  Ascertain  what  happens  when  HC1  is  heated  with  such  sub- 
stances as  manganese  dioxide,  Mn02;  red  lead,  Pb304;  or  potassium 
bichromate,  K2Cr207.     Be  cautious  in  inhaling  the  vapors. 

63.1  In  the  work  with  chlorine  extreme  care  must  be  exercised  not  to 
allow  the  fumes  to  escape  into  the  room.  Keep  all  receivers  well  covered 
with  paper.  If  the  gas  is  accidentally  inhaled,  the  antidote  is  vapor  of  alco- 
hol inhaled  from  a  handkerchief;  or  ammonia. 

Into  a  flask  fitted  with  both  a  delivery  and  a  safety  tube  put  about 
20  g.  of  manganese  dioxide,  Mn02.  Pour  upon  it  enough  ordinary 

1  Unless  there  is  excellent  draught  in  the  laboratory,  the  instructor  is  to 
perform  experiments  63,  64,  65,  and  66  in  the  lecture  room  only. 


26  LABORATORY   EXERCISES. 

concentrated  hydrochloric  acid  to  cover  it  completely.  The  delivery 
tube  should  be  bent  downward  and  reach  nearly  to  the  bottom  of  the 
receiver,  and  should  pass  through  a  hole  in  a  paper  cover  for  the 
receiver.  Shake  the  contents  of  the  flask  well  together  and  apply  very 
gentle  heat.  Fill  several  bottles  with  the  gas.  You  can  tell  when 
they  are  full  by  the  color. 

Note  the  specific  gravity  of  the  gas. 

After  the  necessary  amount  of  gas  has  been  collected  in  the  receiver, 
let  the  gas  pass  into  water  for  a  few  minutes  and  note  if  it  is  soluble. 

Complete  the  equation :  Mn02  +  HC1  =  MnCl2  +  2  H20  +  ? 

64.  Into  a  jar  of  chlorine,  Cl,  thrust  a  burning  taper  or  a  bit  of 
flaming  paper.     Is  the  gas  combustible  ? 

65.  Into  one  of  the  vessels  containing  chlorine  introduce  a  little 
(as  much  as  you  can  put  on  a  ten-cent  piece)  finely  powdered  an- 
timony ;  or  heat  a  small  piece  of  copper  foil  and  introduce  it  into  the 
jar.    What  takes  place  ?    Equation. 

In  what  respects  is  this  experiment  like  the  one  in  which  iron  was 
burned  in  oxygen  ? 

66.  Into  a  vessel  put  a  piece  of  paper  with  some  writing  on  it ; 
some  flowers,  and  some  pieces  of  colored  calico  which  you  have 
moistened,  and  also  pieces  of  written  and  printed  paper. 

What  takes  place  ? 

Into  a  fourth  vessel  put  a  dry  piece  of  the  same  material. 
What  difference  is  there  in  the  action  of  the  chlorine  on  the  dry  and 
on  the  moist  material  ? 

Printer's  ink  is  made  of  lampblack  (carbon)  and  is  not  bleached. 
How  could  you  distinguish  between  organic  and  inorganic  colors  ? 
State  the  theory  of  the  chemistry  of  bleaching. 

67.  Put  into  a  small    beaker   5   g.   bleaching  powder,   CaCl2  + 
Ca(C10)2 ;  set  this  in  a  large  beaker,  and  hang  in  the  latter  the  sub- 
stance to  be  bleached.     Cover  the  large  one  with  pasteboard,  through 
which  passes  a  thistle  tube  into  the  smaller. 

Pour  through  the  thistle  tube  5  cc.  dilute  H2S04.  Add  more  if 
needed. 

Explain  the  action. 

68.  Try  the  effect  of  the  chlorine  water  made  in  Exp.  63  on  a  solu- 
tion of  litmus,  indigo,  or  cochineal.     Also  to  solutions  of  silver  nitrate, 
AgN03,  and  lead  nitrate,  Pb(N03)2,  in  separate  test  tubes,  add  a  little 
of  the  chlorine  water. 

Describe  the  results  and  write  equations. 


LABORATORY   EXERCISES.  27 

How  does  the  action  in  this  case  compare  with  that  of  HC1  on  the 
same  solutions  ?  (See  Exp.  61.) 

69.  Drop  into  a  test  tube  3  or  4  crystals  of  KC103.     Add  a  few 
drops  of  HC1 ;  hold  in  the  flame  for  a  minute,  and  when  action  begins 
add   5  or  10  cc.  H20.     Cautiously  take  the  odor.     What  has  been 
liberated? 

To  2  cc.  indigo  solution  in  a  test  tube  add  a  little  Cl  water. 

Is  the  color  discharged  ? 

To  2  cc.  cochineal  solution  add  a  little  Cl  water.  Is  the  solution 
bleached?  Try  also  litmus  solution. 

To  2  cc.  K2Cr207  solution  add  a  little  Cl  water.  Is  this  bleached? 
K2Cr207  is  a  mineral  pigment ;  cochineal  is  of  animal  origin.  Explain 
the  results. 

This  is  the  ordinary  method  of  making  chlorine  water  for  laboratory 
uses. 

Compare  chlorine  with  hydrogen  and  oxygen. 

EXERCISE  XIX.    BROMINE  AND  IODINE. 

§§  104  to  117. 

70.  From  the  instructor  receive  into  a  flask  or  bottle  of  1  or  2  1. 
capacity  3  or  4  drops  of  bromine,  Br.     Cover  the  bottle  loosely  and 
leave  it  standing.     Immerse  a  piece  of  moist  litmus  paper  in  the  gas. 
What  is  the  effect  ?    Pour  the  liquid  into  a  beaker  of  water,  and  note 
the  specific  gravity. 

71.  Warm  gently  a  few  crystals  of  KBr  with  0.2  g.  Mn02  and  1  cc. 
H2S04  in  a  test  tube  and  observe  the  vapor.     (Under  the  hood.) 

KBr  +  Mn02  +  H2S04  =  NaHS04  +  MnS04  +  H20  +  ? 

Could  Cl  be  made  in  a  similar  manner? 

Illustrate  by  writing  an  equation. 

Test  the  remaining  liquid  with  litmus  paper. 

72.  Make  a  solution  of  a  few  crystals  of  potassium  bromide,  KBr, 
in  3  or  4  cc.  of  water.     Add  a  drop  of  chlorine  water  (see  Exp.  69). 

Is  Br  set  free?    Write  the  equation. 

Add  to  the  above  solution  two  drops  of  carbon  bisulphide  and  shake 
the  tube.  What  effect  is  produced  ? 

How  could  free  bromine  be  detected  in  a  solution  ? 

Why  does  carbon  bisulphide  not  become  colored  when  shaken  in  a 
simple  solution  of  KBr  ? 


28  LABORATORY   EXERCISES. 

EXERCISE  XX.      IODINE  (continued). 
§§  117  to  126. 

73.  Examine  a  small  crystal  of  iodine,  I.    How  does  it  act  upon  the 
fingers  ?     Is  it  soluble  in  water  ?     Try  alcohol. 

What  is  tincture  of  iodine  ? 

74.  Hold  a  dry  test  tube  in  the  gas  lamp  by  means  of  the  wooden 
nippers,  and  warm  it  along  its  entire  length,  so  far  as  this  is  prac- 
ticable.    Drop  into  the  hot  tube  a  small  fragment  of  iodine. 

What  is  the  color  of  the  vapor  V 

75.  Prepare  a  quantity  of  thin  starch  paste  by  boiling  30  cc.  of 
water  in  a  porcelain  dish  and  stirring  into  it  0.5  g.  of  starch  which 
has  previously  been  reduced  to  the  consistency  of  cream  by  rubbing 
it  in  a  mortar  with  a  few  drops  of  water.     Note  the  change  in  the 
starch. 

76.  Pour  3  or  4  drops  of  the  paste  into  10  cc.  of  water  in  a  test 
tube,  and  shake  the  mixture  so  that  the  paste  may  be  equally  diffused 
through  the  water;  then  add  a  drop  of  an  aqueous  solution  of  iodine. 
Heat  the  solution  gently  until  the  color  disappears,  and  allow  it  to 
cool  again.     This  action  affords  a  delicate  test  for  iodine  when  not  in 
combination. 

77.  Dip  a  strip  of  white  paper  in  the  starch  paste  and  suspend  it, 
while  still  moist,  in  a  large  bottle,  into  the  bottom  of  which  2  or  3 
crystals  of  iodine  have  been  thrown. 

What  does  the  experiment  show  ? 

78.  To  a  portion  of  the  starch  paste  made  in  Exp.  75,  add  a  few 
drops  of  potassium  iodide  solution.     Into  the  paste  thus  prepared  dip 
strips  of  filter  paper.     This  is  "  iodo-starch  paper." 

79.  Repeat  Exp.  76,  using  bromine  water  instead  of  an  aqueous 
solution  of  iodine. 

80.  Put  1  g.  Mn02  in  a  test  tube,  pour  upon  it  1  cc.  of  HC1,  and 
warm  gently.     Now  hold  a  piece  of  iodo-starch  paper  over  the  tube 
and  notice  the  result.     Explain  the  action.     This  affords  a  test  for 
chlorine. 

81.  Dissolve  a  few  crystals  of  potassium  iodide,  KI,  in  3  or  4  cc.  of 
water.     Add  to  this  a  little  starch  paste.   Does  it  stain  as  the  element 
did?     Why  not? 

Now  add  some  chlorine  water.     What  effect  ?    Write  equation. 
Which  has  the  stronger  chemism,  chlorine  or  iodine? 


LABORATORY   EXERCISES.  29 

82.  How  could  you  detect  the  presence  of  starch  ? 

How  could  you  distinguish  between  the  three  elements,  Cl,  Br,  and 
I  ?  Classify  these  elements  in  accordance  with  their  properties. 

EXERCISE  XXI.    SULPHUR. 

§§  126  to  136. 

83.  Place  1  g.  of  sulphur  in  a  dry  test  tube  and  pour  upon  it  5  cc. 
of  carbon  bisulphide,  cork  tightly,  and  shake  for  a  few  moments.   Car- 
bon bisulphide  is  volatile  and  very  inflammable.   Have  no  lights  near  by. 
Pour  a  little  of  the  clear  solution  upon  a  watch  glass  and  allow  the 
carbon  bisulphide  to  evaporate  under  the  hood. 

Examine  the  residue  obtained,  and  note  the  shape  of  the  crystals. 

84.  Put  10  g.  S  into  a  test  tube  and  slowly  melt  it.     Notice  the 
yellow  color,  and  see  that  the  liquid  is  very  thin.     It  is  now  somewhat 
above  100°. 

Heat  it  more  strongly  till  it  becomes  black.  It  is  now  very  thick 
and  cannot  be  poured  (200°). 

Apply  more  heat  till  it  grows  thin  again  (300°) . 

Heat  to  boiling  (over  400°) ;  note  the  color  of  the  vapor,  and  any 
sublimate  on  the  test  tube. 

Pour  the  S  into  water.     Knead  it,  and  note  its  elasticity. 

See  whether  it  afterward  changes. 

What  is  the  product  of  burning  sulphur  ?    Write  the  equation. 

85.  Place  in  a  test  tube  a  little  litharge,  PbO,  and  pour  upon  it  a 
few  drops   of  the  H2S   solution.     The  brown   litharge  changes  to 
black.     Lead  oxide  changes  to  lead  sulphide.     Give  the  equation. 

Similar  reactions  take  place  with  other  metallic  oxides  when  treated 
with  H2S. 

86.  Mix  sulphur  and  iron  filings  in  the  proportion  of  3.2  g.  of  the 
former  and  5.6  g.  of  the  latter,  and  heat  the  mixture  in  a  test  tube. 
After  the  mass  glows,  allow  it  to  cool,  and  remove  it  from  the  tube. 
Examine  the  mass  carefully.     Is  it  sulphur  that  remains?    Is  it  iron? 
Try  iron  and  some  of  this  mass  in  dilute  HC1.     What  is  the  mass  ? 
What  does  the  experiment  illustrate? 

Into  5  cc.  of  a  solution  of  lead  nitrate,  Pb(N03)2  pour  a  few  drops 
of  the  H2S  solution.  What  takes  place?  Write  the  equation. 

Add  a  little  H2S  solution  to  each  of  the  following  solutions,  describe 
results,  and  write  the  necessary  equations,  HgCl2,  CuS04,  CdCl2,  BaCl2, 
CaCl2. 


30  LABORATORY  EXERCISES. 

EXERCISE  XXII.    HYDROGEN  SULPHIDE. 
§§  136  to  151. 

87.  Put  into  a  large  side-neck  test  tube  5  g.  ferrous  sulphide,  10  cc. 
water,  and  5  cc.  HC1.     Equation. 

Adjust  a  delivery  tube,  and  pass  the  gas  for  a  minute  or  two  into 
5  cc.  HgO.  Have  the  bearings  tight. 

See  whether  this  solution  is  acid,  alkaline,  or  neutral.  Use  both 
colors  of  litmus. 

Put  a  drop  of  the  H2S  solution  on  Ag  and  Cu  coins.     Reactions. 

Put  a  drop  of  Pb(C2H302)2  solution  on  paper,  and  hold  it  in  the 
vapor  of  H2S.  This  is  the  characteristic  test  for  H2S. 

Mix  equal  parts  of  CuCl2  and  BaCl2,  add  H2S,  shake  and  filter. 
What  is  in  the  filtrate  ?  What  is  on  the  filter  ? 

88.  Referring  to  Exp.  87,  state  how  you  could  test  for  a  sulphide. 
Try  Na2S  or  CaS.     Write  equation. 

EXERCISE  XXIII.    SULPHUR  DIOXIDE. 
§§  151  to  156. 

89.  Light  a  piece  of  sulphur  in  a  deflagrating  spoon,  and  suspend 
the  latter  in  a  cylinder  full  of  air.    Carefully  observe  the  odor. 

Immerse  a  lighted  taper  in  the  gas  obtained. 

90.  Place  a  piece  of  copper  foil  in  a  test  tube,  cover  with  concen- 
trated sulphuric  acid,  and  gently  warm.     Observe  the  odor.     Place  over 
the  mouth  of  the  tube  some  flowers.     What  change  takes  place  in  the 
flowers  ? 

Does  sulphur  dioxide  act  in  the  same  way  that  chlorine  does  ? 

91.  Put  eight  or  ten  pieces  of  sheet  copper,  1  to  2  inches  long  and 
about  £  an  inch  wide,  into  a  flask.     Pour  15  to  20  cc.  concentrated 
sulphuric  acid  upon  it.     Heat  gently.     The  moment  the  gas  begins  to 
come  off,  lower  the  flame  and  keep  it  at  such  a  height  that  the  evolu- 
tion is  regular  and  not  too  active.    Pass  some  of  the  gas  into  a  bottle 
containing  water. 

Is  it  soluble  in  water  ? 

Collect  a  vessel  full  by  displacement  of  air.  (It  is  more  than  twice 
as  heavy  as  air.) 

See  whether  the  gas  will  burn  or  support  combustion. 

Is  the  gas  colored  ?  Is  it  transparent  ?  Has  it  any  odor  ?  Does  it 
burn? 


LABORATORY   EXERCISES.  31 

92.  Charge  a  bottle,  of  the  capacity  of  a  liter  or  more,  with  sulphu- 
rous acid  by  burning  in  it  a  bit  of  sulphur.     Fasten  a  shaving,  or, 
better,  a  tuft  of  gun  cotton,  upon  a  glass  rod  or  tube  bent  at  one  end 
in  the  form  of  a  hook ;  wet  the  shaving  in  concentrated  nitric  acid, 
and  hang  it  in  the  bottle  of  sulphurous  acid.     Interpret  what  you 
observe. 

Pour  a  little  BaCl2  solution  into  the  bottle  before  beginning  the 
experiment,  and  notice  its  condition  after  shaking  the  bottle. 
Write  the  equation. 

EXERCISE  XXIV.    SULPHURIC  ACID. 

§§  156  to  164. 

93.  Place  in  a  beaker  20  cc.  water ;  pour  gradually  into  the  water 
about  the  same  volume  of  concentrated  sulphuric  acid,  stirring  the 
mixture.     Note  the  change  of  temperature. 

Save  the  dilute  acid. 

94.  Put  one  drop  of  strong  H2S04,  and  one  from  that  just  made,  on 
writing  paper,  and  evaporate  them  high  over  a  flame,  so  as  not  to 
burn  the  paper.     When  it  is  dry,  examine. 

95.  Put  2  cc.  of  strong  H2S04  into  a  test  tube,  and  dip  into  it  a 
splinter.     Wood  and  paper  are  mostly  cellulose,  C18(H2O)15.     Explain 
the  charring. 

96.  To  2  cc.  sugar  solution,  C12(H20)n,  add  2  cc.  H2S04,  and  explain. 
Cover  a  fragment  of  starch,  C6(H20)5,  with  H2S04  in  a  test  tube ;  boil 

till  it  begins  to  blacken.     Explain. 

97.  Pour  a  few  drops  of  H2S04  into  a  test  tube,  and  dilute  with  20  cc. 
of  water.     Add  a  little  BaCl2  solution,  and  observe  the  effect.     This  is 
the  test  for  H2S04. 

98.  Dissolve  a  crystal  of  Pb  (N0a)2  in  water  in  a  test  tube,  and  add 
a  few  drops  of  dilute  H2S04.     Note  the  action.     This  is  another  test 
for  H2S04. 

EXERCISE  XXV.    VOLUMETRIC  COMPOSITION  OF  HYDROCHLORIC 

ACID. 

99.  For  the  Instructor.    Arrange  an  apparatus  for  the  generation  of 
hydrochloric  acid  gas  as  follows :   On  a  ring  of  the  stand  place  an 
ordinary  500  cc.  Erlenmeyer  flask,  fitted  by  a  two-hole  rubber  stopper, 
with  a  glass  stopcock  funnel  passing  nearly  to  the  bottom  of  the  flask, 
and  a  glass  tube  bent  at  a  right  angle,  which  tube  connects  with  a 


32  LABORATORY   EXERCISES. 

similar  tube  leading  to  the  bottom  of  a  500  cc.  Erlenmeyer  filter 
flask  with  a  side  neck,  to  which  is  connected  about  a  foot  of  rubber 
tubing  which  can  be  closed  air-tight  by  means  of  a  stout  pinchcock. 
Place  concentrated  sulphuric  acid  in  the  funnel  (stopcock  closed), 
and  common  salt  in  the  generating  flask.  Allow  the  acid  to  drop 
slowly  on  the  salt,  thus  generating  HC1.  Warm  the  generating  flask 
occasionally.  Completely  fill  the  filtering  flask  with  the  gas  by  down- 
ward displacement.  Now  put  into  the  filtering  flask  containing  the 
HC1  3  cc.  of  magnesium  powder,  and  stopper  the  flask  air-tight,  plac- 
ing the  free  end  of  the  rubber  tube  attached  to  the  side  neck  in  a 
beaker  of  water.  Very  cautiously  open  the  pinchcock  on  the  tube  and 
allow  the  water  to  pass  into  the  flask  to  about  half  its  capacity.  Close 
the  pinchcock  again  (make  sure  it  is  tight)  and  allow  to  stand  12 
hours  or  more,  to  complete  reaction.  Invert  the  flask  to  bring  the 
water  into  the  neck,  open  the  pinchcock,  and  make  the  level  of  the 
water  the  same  in  the  flask  and  the  beaker.  Why? 

Set  the  flask  upright  and  measure  the  water,  and  also  ascertain 
the  total  contents  of  the  flask,  making  the  necessary  corrections. 
Remembering  that  the  flask  was  at  first  full  of  HC1,  which  was  then 
absorbed  by  the  water  and  finally  decomposed  by  the  magnesium, 
as  follows :  — 

2HCl+Mg  =  MgCl2+2H. 

Calculate  the  relative  proportion  of  chlorine  and  hydrogen  in  the 
original  HC1  by  volume. 

100.  Repeat  Exp.  99,  using  2  volumes  of  hydrogen  to  1  of  oxygen. 
How  many  volumes  of  gas  remains  after  the  combination  ?  What  is 
this  gas  ?  What  was  the  condensation  ratio  in  this  case  ?  Why  was 
there  not  a  condensation  in  the  previous  experiment  ? 

The  resulting  volume  is  a  constant  one  for  the  union  of  gases. 
Make  a  table  of  the  nitrogen  oxides,  showing  the  condensation  in  each 
case. 

From  the  above  what  may  be  learned  of  how  the  space  occupied  by 
the  compound  molecule  compares  with  that  occupied  with  the  unit 
volume  ? 

This  double  volume  is  often  called  the  product  volume  of  a  compound 
gas. 

EXERCISE  XXVI.    MOLECULAR  WEIGHTS. 

For  the  argument  on  "  Molecular  Condition  of  Gases  "  the  student  is  referred 
to  Storer  and  Lindsay's  "  Manual  of  Chemistry,"  §§  168  to  173. 


LABORATORY   EXERCISES.  33 


DETERMINATION  OF  MOLECULAR  WEIGHTS. 

Accepting  Avogadro's  Law,  and  the  dependent  fact  that  the  vapor 
density  equals  one  half  the  molecular  weight,  it  becomes  possible  to  calcu- 
late the  molecular  weight  of  any  substance  which  is  naturally  a  gas  or 
which  can  be  easily  vaporized.  This  is  done  by  the  physical  method. 

101.  Fit  as  large  a  flask  as  will  ride  conveniently  on  the  laboratory 
balance  with  a  one-hole  rubber  stopper,  through  which  passes  a  tightly 
fitting  glass  tube  projecting  an  inch  above  the  stopper,  and  reaching 
to  the  bottom  of  the  flask.     Close  the  outer  end  of  the  glass  tube  by 
means  of  rubber  tube  and  pinchcock.     Be  sure  that  all  joints  are  per- 
fectly tight.     Fill  the  flask  to  the  pinchcock  with  water,  and  ascertain 
by  measurement  the  volume  of  the  flask.     Clean  and  dry  the  flask,  and 
loosen  the  stopper,  but  do  not  remove  it.     Now  counterpoise  the  flask 
on  the  balance,  or  ascertain  the  weight,  and  record  the  same  in  your 
notebook.     Ascertain  both  the  temperature  at  the  balance  and  the 
barometer  reading.     Now  calculate  according  to  Exp.  22  the  weight 
of  air  in  the  open  flask  (1  cc.  of  air  at  standard  conditions  weigh 
.001293  g.). 

Generate  oxygen,  passing  it  first  through  water  and  then  through 
concentrated  sulphuric  acid  contained  in  wash  bottles.  Allow  the  gas 
to  pass  through  the  tube  of  the  weighing  flask  till  a  glowing  match 
held  near  the  loose  stopper  shows  the  flask  to  be  full  of  the  gas. 
Insert  the  stopper,  and  close  the  pinchcock.  Allow  any  excess  to  es- 
cape by  opening  for  an  instant  only  the  pinchcock.  Read  the  tempera- 
ture, the  barometer,  and  find  the  weight  of  the  flask.  Repeat  the 
filling  of  the  flask  till  there  is  no  further  gain  in  weight.  As  soon 
as  a  constant  weight  has  been  reached,  again  take  readings  of  tem- 
perature and  pressure,  using  the  latter  readings  for  the  computation. 
The  gain  in  weight,  plus  the  weight  of  the  air  the  flask  held,  is  the 
weight  of  the  oxygen  at  the  temperature  and  pressure  observed.  Cal- 
culate the  weight  of  the  same  volume  at  standard  conditions.  The 
weight  varies  directly  as  the  pressure  and  inversely  as  the  absolute 
temperature.  Why?  Calculate  the  density  as  referred  to  hydrogen. 
From  this  calculate  the  molecular  weight  of  oxygen. 

EXERCISE  XXVII. 

102.  Many  substances  cannot  be  vaporized,  hence  it  is  impossible 
to  use  the  physical  method  above  illustrated.     Recourse  in  such  cases 
is  had  to  chemical  methods.    Assuming  that  the  molecular  weight  of 

LAB.  EX.  —  3 


34  LABORATORY   EXERCISES. 

oxygen  has  been  ascertained  to  be  32,  we-  can  now  proceed  to  deter- 
mine the  molecular  weight  of  a  substance  in  which  oxygen  is  a  com- 
ponent, as,  for  instance,  potassium  chlorate,  KC10r 

For  this  experiment  the  perfectly  pure  and  dry  salt  must  be  used. 
Weigh  into  a  small  porcelain  crucible  about  2  g.  of  the  pulverized  dry 
salt.  The  weight  must  be  accurately  known,  and  the  crucible  should  be 
provided  with  a  cover.  By  means  of  a  Bunsen  flame  heat  the  chlorate 
in  the  crucible  gently  at  first,  taking  care  all  the  time  to  avoid  spat- 
tering and  foaming.  Continue  the  heating  till  the  mass  becomes 
solid  after  melting.  Apply  heat  by  means  of  the  blast  lamp  till  the 
mass  again  melts,  after  which  remove  to  a  desiccator,  let  cool,  and 
weigh.  Repeat  the  heating  with  the  blast  lamp,  cooling  and  weighing 
till  there  is  no  further  loss.  Oxygen  has  been  driven  off  (see  Exp.  13). 
Calculate  the  loss  due  to  escape  of  oxygen. 

In  this  experiment  the  chlorate  gives  up  all  its  oxygen,  and  the 
chemical  changes  in  which  it  has  entered  shows  it  to  have  three  atoms 
of  oxygen,  hence  a  molecule  of  the  salt  has  given  enough  oxygen  to 
make  one  and  a  half  molecules  of  oxygen  gas.  Form  the  propor- 
tion :  — 
wt.  of  oxygen  given  off :  wt.  in  grams  of  chlorate  taken : : 

mol.  wt.  of  1^  molecules  of  oxygen  :  x  (=  mol.  wt.  of  KClOs). 

EXERCISE  XXVIII.    PHOSPHORUS. 
§§  182  to  194. 

In  handling  phosphorus  extreme  care  must  be  exercised.  Do  not  touch 
it  with  the  Jingers.  Always  cut  it  under  water,  holding  it  with  pincers. 
Always  have  water  at  hand  to  extinguish  it  if  premature  ignition  takes 
place. 

103.  Put  a  piece  of  P  as  large  as  a  grain  of  wheat  in  a  test  tube 
and  pour  on  it  immediately  1  cc.  of  carbon  bisulphide,  CS2.  Does  the 
P  dissolve  ? 

Pour  a  little  of  the  liquid  on  a  piece  of  filter  paper,  leaving  enough 
of  the  paper  dry  to  hold  it  by ;  wave  it  back  and  forth  till  the  CS2  has 
evaporated.  The  P  ignites  spontaneously.  Does  the  paper  burn? 
Why  not? 

Why  does  not  the  P  on  a  match  burn  spontaneously  ? 

What  is  the  product  formed  when  P  burns  ?    Write  equation. 

Why  is  P  kept  under  water  ?  Why  does  it  not  combine  with  the  0 
of  the  water  ? 


LABORATORY   EXERCISES.  35 

104.  In  an  evaporating  dish  place  a  grain  of  P  as  large  as  a  grain 
of  rice,  cover  it  with  a  beaker,  and  ignite  it  by  means  of  a  warm 
glass  rod.     Note  the  white,  flaky  product  formed.     What  is  it  ? 

Breathe  on  it.     What  result  ? 

Wash  the  inside  of  the  beaker  with  a  little  water  and  test  the 
water  with  litmus  paper. 

In  this  experiment  metaphosphorous  acid  is  formed.  Write  the 
equation. 

105.  In  an  evaporating  dish  put  about  2  g.  of  "  glacial "  phosphoric 
acid,  HP03,  add  about  3  oc.  of  water.     After  a  little  of  the  acid  has 
dissolved  pour  the  liquid  into  a  test  tube  and  add  a  solution  of  silver 
nitrate,  AgN03.     White  precipitate.     Write  equation. 

To  the  HP03  remaining  in  the  dish  add  about  30  cc.  water.  Boil  a 
few  minutes.  Test  a  little  of  the  liquid  with  AgN03 ;  then  add  1  or  2 
drops  of  water. 

Complete  the  equation  :     HP03  +  H20  =  H4P207 

Write  equation  for  the  action  of  AgN03  on  H4P207. 

Now  fill  the  dish  with  water  and  boil  for  about  30  minutes. 

Test  again  with  AgN03,  and  the  smallest  drop  of  NH4OH. 

Are  the  precipitates  the  same  in  each  case  ? 

Complete  the  equations  :  — 

H4P207  +  H20  =  H3P04 

H3P04  +  AgN03  =  ? 

* 

106.  To  10  cc.  of  Na2HP04  solution  add  BaCl2  in  excess.     Filter  the 
mixture,  and  test  the  solubility  of  the  precipitate  in  ammonia  water 
and  in  dilute  hydrochloric  acid. 

107.  (I)  To  2  cc.  of  Na2HP04  add  10  cc.  of  water,  2  cc.  of  NH4C1, 
and  1  cc.  of  ammonia  water.     Now  add  slowly  3  cc.  of  a  mixed  solu- 
tion containing  MgS04  and  NH4C1. 

Observe  carefully  under  the  microscope  the  character  of  the  precipi- 
tate formed. 

108.  (II)  To  5  cc.  of  Na2HP04,  add  AgN03  solution  in  slight  excess. 
Note  the  color  of  the  precipitate.     Filter  and  test  its  solubility  in 
ammonia  water,   dilute  nitric  acid,   and  dilute  hydrochloric    acid. 
Save  the  silver  residue. 

109.  (Ill)  To  a  few  cc.  of  Na2HP04  add  an  equal  volume  of  a 
solution  of  ammonium   molybdate  in  nitric  acid.     The  last  three 
experiments  are  the  usual  tests  for  phosphoric  acid  and  phosphates. 


36  LABORATORY   EXERCISES. 

110.  Boil  a  little  bone  ash  in  water  in  a  tube  and  filter  the  liquid. 
Test  the  liquid  for  phosphates  as  in  Exp.  107.     The  bone  ash  con- 
sists of  tricalcium  phosphate,  Ca3(P04)2,  which  is  nearly  insoluble, 
so  there  will  be  very  little,  if  any,  trace  of  phosphates  in  the  liquid. 

Now  place  a  little  bone  ash  (or  bone  meal)  in  a  test  tube  and  add 
two  or  three  drops  of  sulphuric  acid ;  warm  the  mixture.  After  the 
tube  becomes  cool,  add  water  and  shake  the  mixture ;  filter.  Test 
the  filtrate  as  above.  What  has  been  the  effect  of  the  sulphuric  acid 
on  the  bone  ash?  How  do  you  know? 

The  change  is  represented  as  follows :  — 

Ca3(P04)2  +  2  H2S04  =  CaH4(P04)2  +  ? 

This  experiment  illustrates  the  manufacture  of  superphosphates 
for  fertilizing  purposes. 

EXERCISE  XXIX.     SILICA  AND  CARBON. 
§§  215  to  229. 

111.  To  a  concentrated  solution  of  "  water  glass  "  in  an  evaporating 
dish  add  enough  concentrated  hydrochloric  acid,  HC1,  to  render  the 
solution  acid.     A  jellylike  mass  of  silicic  acid,  H4Si04,  will  separate. 
Evaporate  the  contents  of  the  dish  to  dryness,  slowly,  and  then  heat 
the  residue  gently  over  a  lamp.     After  allowing  to  cool  add  water. 
Examine  the  contents  of  the  dish  for  a  fine  white  powder,  Si02.    Give 
the  reaction  representing  the  change  from  H4Si04  to  Si02. 

112.  Put  into  a  tube  about  12  cm.  long,  enough  soft  coal  to  fill  it 
about  one  third  full.1    Fit  to  the  tube  a  delivery  tube,  and  support 
the  apparatus  on  the  ring  stand.     Heat  the  coal  and  collect  the  gas. 
Test  the  gas  with  a  flame.     It  consists  of  a  mixture  of  carbon  com- 
pounds.    It  is  illuminating  gas.     It  contains  NH3. 

How  could  this  be  removed  ? 

Remove  the  coke  from  the  tube  and  examine  it.  See  if  it  will  burn, 
and  note  in  what  manner. 

113.  Repeat  Exp.  112,  using  wood  shavings  or  saw  dust.    Collect  the 
gas  and  test  it.     After  driving  off  the  gas,  remove  the  end  of  the  tube 
from  the  water,  plug  it  to  prevent  air  entering,  and  allow  the  appa- 
ratus to  cool.     Finally  remove  the  contents  of  the  tube  and  examine. 
How  does  the  charcoal  burn  ? 

1  A  clay  pipe  whose  bowl  is  filled  with  the  coal  and  sealed  with  plaster  of 
Paris  answers  well. 


LABORATORY  EXERCISES.  37 

114.  Over  a  burning  candle  hold  the  bottom  of  an  evaporating  dish 
or  glass  plate.     Note  the  collection  of  lampblack  (carbon). 

Explain  why  the  smoke  accumulates. 

115.  Mix  on  paper  and  put  into  a  tube  10  parts  CuO  to  1  part  C 
(powdered  charcoal),  by  weight  5  g.  in  all.     The  tube  should  not  be 
over  one  third  full.     Pass  the  gas  into  limewater  contained  in  a  test 
tube.    What  result?    Carbon  dioxide  turns  limewater  milky. 

What  is  the  appearance  of  the  substance  left  in  the  tube?  Does  it 
suggest  the  metal  copper,  Cu  ?  Treat  a  little  with  concentrated  nitric 
acid,  HN03.  What  should  take  place  if  the  substance  is  metallic 
copper  ? 

What  does  take  place  ?    Write  equation. 

What  is  the  reaction  which  takes  place  between  the  copper  oxide 
and  the  charcoal?  Write  equation. 

Compare  the  action  of  hydrogen  with  that  of  carbon  on  copper 
oxide.  In  what  respects  are  they  alike,  and  in  what  respects  do  they 
differ? 

EXERCISE  XXX.    CARBON    (continued). 
§§  229  to  243. 

116.  Prepare  a  solution  of  H2S  in  a  receiver  with  20  cc.  H20. 
Notice  the  odor. 

Put  into  the  receiver  5  g.  powdered  charcoal,  and  shake  the  mixture 
well.  Pour  the  whole  on  a  filter,  collect  the  filtrate  in  a  clean  receiver, 
and  see  whether  any  odor  remains.  If  so,  use  more  charcoal  and 
filter  again. 

117.  Make  a  filter  of  boneblack  by  fitting  a  paper  filter  into  a 
funnel  8  to  10  mm.  (3  to  4  inches)  in  diameter  at  its  mouth.     Half 
fill  this  with  boneblack.     Pour  a  dilute  solution  of  indigo l  through 
the  filter. 

What  effect  does  this  have  on  the  color  of  the  solution? 

Do  the  same  thing  with  a  dilute  solution  of  litmus.  If  the  color  is 
not  completely  removed  by  one  filtering,  filter  the  solution  again. 

Try  also  a  solution  of  potassium  bichromate,  K2Cr207. 

As  indigo  and  litmus  are  organic  coloring  matters,  and  K2Cr207  is 
mineral,  state  any  inference. 

1  Prepared  by  treating  |  g.  of  powdered  indigo  for  some  time  with  4-5 
cc.  of  warm  concentrated  sulphuric  acid  and  diluting  with  a  liter  of  water. 


38  LABORATORY   EXERCISES. 

118.  Mix  4  g.  of  potassium  nitrate  with  2  g.  of  powdered  charcoal. 
Place  the  mixture  upon  an  iron  plate  and  touch  it  with  a  lighted 
stick.    What  does  it  act  like  ? 

EXERCISE  XXXI.    CARBON  DIOXIDE. 
§§  243  to  247. 

119.  Put  a  short  piece  of  candle  on  the  desk  and  invert  over  it  a 
dry,  wide-mouthed  bottle.     What  forms  on  the  inside  of  the  bottle  ? 

Remove  the  bottle  and  pour  into  it  a  small  quantity  of  limewater 
and  shake.  What  is  the  effect  on  the  limewater?  What  is  formed? 

Write  the  equation  for  the  product  of  combustion  and  for  the  action 
of  the  limewater. 

120.  Into  a  beaker  of  limewater  blow  the  breath  through  a  glass 
tube.     Describe  the  result.     Continue  to  blow  till  the  solution  finally 
clears. 

What  is  the  effect  of  'a  saturated  solution  of  C02  on  calcium  car- 
bonate? 

121.  In  a  flask,  arranged  as  for  the  generation  of  hydrogen,  place 
10  or  12  gr.  of  calcium  carbonate  (limestone),  CaC03.     Cover  the  end 
of  the  thistle  tube  with  water,  and  add  by  small  successive  portions 
strong  hydrochloric  acid.     Collect  several  bottles  of  the  gas.     Try 
to  collect  it  by  downward  displacement. 

Thrust  a  lighted  taper  into  a  bottle  of  the  gas.     Note  the  result. 
From  a  large  bottle  of  the  gas  pour  a  quantity  on  a  lighted  candle. 
Is  the  gas  heavier  or  lighter  than  air  ? 

122.  Fill  a  bottle  half  full  of  the  gas.     Cork  under  water  and  shake. 
Lower  the  mouth  of  the  bottle  into  the  water  again,  remove  the  stop- 
per, and  note  if  the  water  rises  on  the  inside.     To  what  is  this  due  ? 

What  is  "  soda  water  "  ?    What  other  gases  will  dissolve  in  water? 

123.  Pass  carbon  dioxide  into  a  solution  of  2  gr.  of  caustic  potash 
(potassium  hydroxide)  until  it  will  absorb  no  more. 

Add  any  dilute  acid  in  2  or  3  cc.  of  water  in  a  test  tube  to  the  solu- 
tion thus  obtained.  What  gas  is  given  off  when  the  acid  is  added? 
How  do  you  know  ? 

Write  the  equations  expressing  the  reactions  which  take  place  on 
passing  the  carbon  dioxide  into  the  caustic  potash  solution,  and  on 
adding  an  acid  to  the  solution. 

124.  Put  about  a  gram  of  sodium  carbonate,  Na2C03,  into  each  of 
four  test  tubes,  and  then  add  to  one  tube  about  4-5  cc.  of  dilute  hydro- 
chloric acid,  to  a  second  add  the  same  quantity  of  dilute  sulphuric 


LABORATORY   EXERCISES.  39 

acid,  to  a  third  the  same  quantity  of  dilute  nitric  acid,  and  to  the 
fourth  the  same  quantity  of  dilute  acetic  acid.  What  takes  place? 
Is  a  gas  given  off  ? 

Pass  it  through  limewater.     What  is  it  ? 

Perform  the  same  experiment  with  small  pieces  of  marble.  What 
gas  is  given  off? 

What  conclusions  can  you  draw  from  these  observations? 

How  can  you  easily  detect  carbon  dioxide  ? 

EXERCISE  XXXII.      FLAME. 

§§  257  to  267. 

125.  Examine  the  structure  of  a  Bunsen  burner  (unscrew  the  top), 
make  a  drawing  to  show  the  orifices,  and  state  what  use  each  sub- 
serves.    Light  the  gas  at  the  base  for  a  minute. 

Replace  the  top,  and  light  the  gas  at  the  top  of  the  burner. 

Hold  the  flame  in  front  of  a  dark  object,  examine  the  parts,  make  a 
drawing,  give  a  brief  description,  and  state  the  color  of  each  part. 

Put  the  flame  in  direct  sunlight,  and  study  the  parts  from  its  shadow, 
to  confirm  your  results. 

Make  a  careful  examination  of  the  parts  and  colors  of  a  candle  flame, 
and  make  a  drawing  to  show  them.  Move  it  slightly  in  the  air  to 
show  the  outer  flame.  This  is  best  seen  in  a  dark  room. 

126.  Light  the  gas  of  a  Bunsen  burner.     Put  a  stick  across  the 
base  of  the  flame  for  an  instant,  and  notice  what  parts  are  burned. 
Make  a  sketch. 

Hold  a  stick  just  above  the  inner  blue  cone  of  the  flame. 

Press  quickly  down  on  the  flame  with  a  paper  (remove  before  it 
burns)  and  notice  the  shape  of  the  charred  part.  Sketch.  Press  down 
on  the  flame  with  a  fine  wire  gauze,  and  observe  by  the  glowing  of  the 
wire  where  the  heat  is  most  intense. 

Test  the  heat  of  the  inner  cone  with  the  end  of  a  platinum,  Pt,  wire. 
Notice  that  it  glows  when  near  the  top,  but  not  elsewhere  in  this  cone. 

Put  one  end  of  a  small  tube  into  the  inner  blue  cone,  and  try  to 
light  the  gas  at  the  other  end. 

From  the  above,  state  what  takes  place  in  each  of  the  two  chief 
parts  of  the  flame. 

127.  Observe  the  light  of  a  Bunsen  flame,  and  its  color. 
Sprinkle  a  very  little  charcoal  dust  in  the  flame,  and  note  any  change 

of  light  or  color. 


40  LABORATORY   EXERCISES. 

Close  the  orifices  at  the  base  of  the  burner,  and  explain  the  change 
of  light. 

Hold  an  evaporating  dish  in  the  upper  part  of  this  closed  flame  for 
a  minute,  and  notice  deposit. 

Now  open  the  orifices  and  persistently  try  to  burn  off  a  little  of  the 
deposit  from  the  evaporating  dish. 

What  is  the  cause  of  light  in  a  flame  ? 

128.  Ignite  the  gas  and  hold  a  fine  wire  gauze  3  or  4  cm.  above  the 
burner.     Why  does  it  not  burn  above  the  wire? 

Extinguish,  then  relight,  the  gas  above  the  gauze.     Result. 

Gradually  lift  the  wire  till  the  gas  will  not  burn. 

Again  light  the  gas  above  the  gauze,  and  hold  another  gauze  above 
the  flame,  so  as  to  confine  it  above  and  below. 

From  this  experiment  define  kindling  point,  and  state  three  condi- 
tions of  combustion. 

129.  Put  a  fragment  of  lead,  Pb,  not  larger  than  a  pea  on  a  piece 
of  charcoal,  slightly  hollowed  out  to  hold  it.     Insert  the  metallic  tube 
in  a  Bunsen  burner,  and  with  a  mouth  blowpipe  direct  the  oxidizing 
flame  strongly  and  steadily  against  the  Pb  for  4  or  5  minutes. 

As  you  stop  blowing,  notice  the  yellow  vapor  that  escapes  from  the 
pellet  of  Pb;  note  also,  as  it  cools,  the  yellow  coating  of  lead  oxide, 
PbO,  on  the  coal.  Write  the  reaction. 

130.  Put  \  gr.  PbO  on  a  piece  of  charcoal.     With  the  blowpipe  di- 
rect the  reducing  flame  steadily  against  it  for  some  time,  or  until  a 
metallic  pellet  is  obtained. 

What  is  it  ?    Equation. 

EXERCISE  XXXIII.    METHANE  AND  ILLUMINATING  GAS. 

131.  Mix  together  2  g.  of  crystallized  sodium  acetate,  4  g.  of  caustic 
soda,  and  8  g.  of  slaked  lime.     Heat  the  mixture  gently  upon  an  iron 
plate,  until  all  the  water  of  crystallization  of  the  acetate  has  been 
expelled,  and  the  mass  has  become  dry  and  friable.     Charge  an  igni- 
tion tube  20  cm.  long  with  the  dry  powder,  heat  it  above  the  lamp, 
and  collect  the  gas  at  the  water  pan.    Marsh  gas  is  evolved  from  the 
mixture,  at  a  temperature  below  redness,  and  a  residue  of  sodium  car- 
bonate is  left  in  the  ignition  tube.     The  purpose  of  the  lime  is  to  ren- 
der the  mass  porous  and  infusible,  or  nearly  infusible,  so  that  the  tube 
may  be  heated  equably.    The  reaction  may  be  represented  as  follows :  — 

NaC2H302  +  NaOH  =  CH,  + 


LABORATORY   EXERCISES.  41 

132.  Fill  a  tall  bottle  of  at  least  one  liter  capacity  with  warm  water, 
invert  it  over  the  water  pan,  and  pass  marsh  gas  into  it,  until  a  little 
more  than  one  third  of  the  water  is  displaced ;  cover  the  bottle  with  a 
thick  towel,  to  exclude  the  light,  and  then  fill  the  rest  of  the  bottle 
with  chlorine.     Cork  the  bottle  tightly,  and  shake  it  vigorously,  in 
order  to  mix  the  gases  together,  keeping  the  bottle  always  covered 
with  the  towel.     Finally,  open  the  bottle,  and  apply  a  light  to  the 
mixture.     Ignition  takes  place,  hydrochloric  acid  is  produced,  while 
the  sides  and  mouth  of  the  bottle  become  coated  with  solid  carbon  in 
the  form  of  lampblack.      The  presence  of  the  acid  may  be  proven 
by  the  smell,  by  its  reaction  with  moistened  blue  litmus  paper,  and  by 
the  white  fumes  which  are  generated  when  a  rod  moistened  with 
ammonia  water  is  brought  in  contact  with  the  escaping  acid  gas. 

133.  Fill  an  ignition  tube  a  third  full  of  bituminous  coal.     Hold  it 
steadily  over  a  lamp  to  heat  it,  meantime  trying  to  ignite  the  escap- 
ing gas.     Note  the  color  of  the  flame,  and  see  whether  any  soot  is 
deposited  on  porcelain  held  in  the  burning  gas.     Break  and  examine 
the  tube  for  a  tarry  residue  arid  for  coal. 

Put  the  coal  on  an  iron  plate,  and  bring  a  Bunsen  flame  in  contact 
with  it,  noting  whether  it  burns  with  a  flame  or  only  glows.  Only 
gas  burns  with  flame. 

Write  no  reactions,  but  state  how  many  and  what  products  you 
observed  in  this  experiment. 

EXERCISE  XXXIV.    ALCOHOL. 

134.  Introduce  20  cc.  of  molasses  into  a  flask  of  200  cc.,  fill  it  with 
water  to  the  neck,  and  put  in  half  a  cake  of  yeast.     Fit  to  this  a 
delivery  tube,  and  pass  the  end  of  it  into  a  test  tube  holding  a  clear 
solution  of  limewater.     Leave  in  a  warm  place  for  two  or  three  days. 
Then  look  for  a  turbidity  in  the  limewater,  and  account  for  it. 

Is  the  liquid  remaining  in  the  flask  sweet?  What  has  become  of 
the  sugar?  This  is  a  fermented  liquor. 

135.  Pour  off  one  half  of  the  fermented  liquor,  and  reserve  it  in  a 
loosely  covered  dish  ;  with  the  remainder  proceed  as  follows :  Support 
the  flask  on  the  iron  lamp  stand,  and  by  means  of  a  delivery  tube  con- 
nect it  with  a  second  flask  capable  of  holding  one  third  of  the  liquid, 
and  placed  on  a  water  bath.     From  this  second  flask  a  delivery  tube 
is  carried  to  a  small  flask  kept  cool  by  immersion  in  cold  water.    Heat 
the  liquid  in  the  largest  flask,  so  that  it  just  boils.    The  vapor  of  alco- 


42  LABORATORY   EXERCISES, 

hoi,  together  with  a  certain  amount  of  steam,  passes  into  the  second 
flask,  which  is  kept  just  below  the  boiling  point  of  water  by  being 
supported  on  the  water  bath  in  which  the  water  barely  boils.  At  this 
temperature  a  considerable  portion  of  the  alcohol,  together  with  some 
water,  passes  over  into  the  third  flask,  where  it  is  condensed.  Con- 
tinue the  operation  until  about  one  third  of  the  liquid  has  passed  out 
of  the  large  flask.  The  liquid  obtained  in  the  third  flask  is  a  dilute 
alcohol ;  the  odor  of  alcohol  is  distinctly  perceptible,  but  the  alcohol 
may  not  be  strong  enough  to  burn.  In  that  case  support  the  third 
flask  on  the  wire  gauze  over  the  lamp,  and  connect  it  by  means  of  a 
delivery  tube  with  another  small  flask,  which  is  kept  cool.  Heat  the 
contents  of  the  flask  gently  until  they  just  boil,  and  transfer  the  first 
teaspoonful  of  the  liquid  which  condenses  in  the  cooled  flask  to  a 
porcelain  dish.  If  the  experiment  has  been  successfully  conducted, 
the  alcohol  thus  obtained  will  be  strong  enough  to  take  fire  if  a  flame 
be  brought  into  contact  with  it. 

Taste  and  smell  the  distillate.     Try  to  ignite  it. 

136.  Put  a  little  of  the  white  of  egg  into  a  beaker ;  cover  it  with 
strong  alcohol  and  note  the  effect.     Strong  alcohol  has  the  same  coagu- 
lating action  on  the  brain  and  on  the  tissues  generally,  when  taken 
into  the  system,  absorbing  water  from  them,  hardening  them,  and 
contracting  them  in  bulk. 

EXERCISE  XXXV.    ETHER. 

The  student  should  never  attempt  to  perform  any  experiment  requiring 
more  than  a  very  minute  quantity  of  ether,  since  it  is  highly  dangerous  to 
work  with  this  substance  on  account  of  its  great  volatility  and  ready 
inflammability. 

137.  Into  a  small  test  tube  put  10  drops  of  ordinary  alcohol  and  as 
much  strong  sulphuric  acid,  and  heat  the  mixture  gently  over  the 
lamp.     Ether  will  be  formed,  and  may  be  recognized  by  its  peculiar 
odor. 

138.  Pour  a  small  quantity  of  ether  into  the  palm  of  the  hand,  and 
observe  the  rapidity  with  which  it  evaporates,  and  also  the  cold  pro- 
duced by  this  evaporation. 

139.  Into  a  tumbler  or  other  very  wide-mouthed  vessel  put  a  few 
drops  of  ether.     Cover  the  vessel  loosely,  and  allow  to  stand  for  a  few 
moments ;  then  bring  a  lighted  match  to  the  mouth  of  the  vessel : 
the  heavy  vapor  of  ether  will  have  displaced  the  air  in  the  vessel,  and 
will  take  fire  at  the  mouth  of  the  vessel  with  a  sudden  flash. 


LABORATORY   EXERCISES.  43. 

140.  Into  a  small  test  tube  put  10  drops  of  ordinary  alcohol,  and 
the  same  amount  of  strong  sulphuric  acid.     Add  a  crystal  of  sodium 
acetate  as  large  as  a  small  pea,  and  heat  the  mixture  gently.    Acetic 
ether,  ethyl  acetate,  is  formed,  and  may  be  recognized  by  its  peculiar 
odor. 

Examine  the  alcoholic  liquid  preserved  from  Exp.  135  by  taste, 
smell,  and  litmus.  The  alcohol  has  changed  to  what  ?  Preserve  the 
liquid  for  subsequent  work. 

EXERCISE  XXXVI.    ACETIC  ACID  AND  ACETATES. 

141.  To  the  acid  liquid  of  the  previous  experiment  or  to  40  or 
50  cc.  of  common  vinegar,  add  powdered  chalk  (calcium  carbonate) 
as  long  as  the   addition  causes  effervescence.     Calcium  acetate  is 
formed  and  remains  dissolved  in  the  liquid.     Filter  the  solution,  and 
evaporate  the  nitrate  to  dryness  at  a  gentle  heat.     The  solid  residue  is 
an  impure  calcium  acetate.    Place  a  portion  of  this  calcium  acetate  in 
a  small  test  tube,  and  heat  gently  with  a  few  drops  of  strong  sulphuric 
acid.    Acetic  acid  will  be  set  free,  and  may  be  recognized  by  its  pecul- 
iar odor.     If  ordinary  vinegar  be  used  in  this  experiment,  it  will  be 
better  to  decolorize  the  solution  of  calcium  acetate  by  mixing  it  with 
powdered  bone  black  before  filtering. 

SOAP. 

142  (a).  Dissolve  15  g.  of  solid  caustic  soda  in  120  cc.  of  water. 
When  the  suspended  impurities  have  settled  to  the  bottom  of  the 
solution,  pour  off  one  half  of  the  clear  liquor  into  a  deep  iron  or 
porcelain  dish  of  at  least  500  cc.  capacity,  add  an  equal  bulk  of  water, 
and  50  g.  of  beef  tallow.  Bring  the  mixture  to  boiling  and  boil  it 
steadily  for  three  quarters  of  an  hour,  supplying  from  time  to  time 
the  water  .lost  by  evaporation;  then  add  the  remainder  of  the  solution 
of  caustic  soda,  and  continue  to  boil  steadily  for  an  hour  or  more, 
allowing  the  liquid  to  become  somewhat  more  concentrated  toward 
the  end  of  that  time ;  then  add  20  g.  of  fine  salt,  boil  for  a  minute  or 
two,  and  allow  the  liquid  to  cool.  A  part  of  the  mass  becomes  solid, 
and  rises  to  the  top ;  it  is  hard  soap. 

The  chemical  action  is  thus  explained:  when  tallow  (glyceryl 
stearate  and  oleate)  is  boiled  with  sodium  hydroxide,  there  is  formed 
sodium  stearate  (and  oleate)  and  glyceryl  hydroxide.  When  common 
salt  is  added,  the  soap  (sodium  stearate  and  oleate),  being  insoluble 


44  LABORATORY   EXERCISES. 

in  the  saline  liquid,  separates  as  a  solid.  The  liquid  remaining  con- 
tains in  solution  the  excess  of  sodium  hydrate  employed,  as  well  as 
the  salt  and  the  glycerine. 

(&).  Soap  may  be  made  more  quickly  by  using  castor  oil  in- 
stead of  beef  tallow.  Mix  100  cc.  of  castor  oil  and  100  cc.  of  caustic 
soda  solution  prepared  as  above,  and  boil  for  30  minutes.  Then  add 
150  cc.  of  water,  bring  to  a  boil,  and  add  20  g.  of  salt.  The  soap  rises 
to  the  top  and  may  be  removed  when  cold.  Castor  oil  is  mainly 
glyceryl  ricinoleate;  the  chemical  change  is  similar  to  that  just 
described. 

143.  Heat  some  of  the  soap  with  soft  water.     A  nearly  clear  solu- 
tion will  be  obtained  if  the  decomposition  of  the  tallow  or  oil  was 
complete.     Add  dilute  hydrochloric  acid  until  the  solution  is  decidedly 
acid.     The  liquid  will  become  turbid,  and  on  standing  will  become 
covered  with  a  layer  of  a  fatty  substance,  which  is  a  mixture  of  stearic 
and  oleic  acids  (or  mainly  ricinoleic  acid  if  castor  oil  was  used).     The 
sodium  chloride  formed  will  be  held  in  solution  by  the  liquid. 

144.  Dissolve  some  of  the  soap  above  made  in  water,  and  render 
the  water  hard  by  adding  10  cc.  of  a  clear  solution  of  calcium  or 
magnesium  sulphate.     Shake  and  note  any  change.     An  insoluble 
soap  has  been  formed. 

f  stearate    1 

Sodium  j  palmitate  t  +  calcium  sulphate  =  ? 
I     oleate     J 

EXERCISE  XXXVII.    SUGAR. 

145.  Heat  cautiously  a  small  quantity  of  white  sugar  in  a  porcelain 
dish  until  it  melts.    Allow  the  pasty  liquid  to  cool  rapidly.    The  product 
is  barley  sugar.    Heat  again  to  a  higher  temperature,  but  not  too  high  ; 
the  sugar  turns  brown,  froths,  and  gives  off  pungent  vapors,  and  there 
remains  a  dark  brown  mass.     This  is  caramel. 

146.  Into  a  flask  of  250  cc.  capacity  introduce  100  cc.  of  water.    Add 
1  cc.  of  strong  sulphuric  acid,  and  heat  the  mixture  to  boiling.     In  a 
porcelain  mortar  rub  10  g.  of  starch  with  enough  water  to  make  a 
cream,  and  pour  the  mixture,  little  by  little,  into  the  boiling  liquid, 
taking  care  not  to  interrupt  the  boiling.     The  starch  dissolves  without 
forming  a  paste.     Boil  for  three  or  four  hours,  replacing  from  time 
to  time  the  water  lost  by  evaporation,  and  then  add  powdered  chalk 
(calcium  carbonate)  until  the  liquid  is  no  longer  acid.    When  the 


LABORATORY   EXERCISES.  45 

mixture  has  become  cold,  filter  off  the  insoluble  calcium  sulphate 
formed  by  the  action  of  the  sulphuric  acid  on  the  calcium  carbonate, 
and  evaporate  the  solution  at  a  gentle  heat  to  a  sirupy  consistency. 
The  solution  contains  dextrose,  which,  on  long  standing,  may  separate 
from  the  liquid  in  crystals. 

EXERCISE  XXXVHI.    SOLUBILITY. 

147.  Solubility  in  water.     Weigh  three  portions  of  1  gr.  each  of 
powdered  Na2S04,  CaS04,  and  PbS04.     Take  three  test  tubes  and  place 
in  each  10  cc.  of  water.     Into  one  of  the  tubes  pour  one  portion  of 
Na2S04  and  shake,  and,  if  this  dissolves,  add  the  second ;  and,  if  the 
second  dissolves,  add  the  third.     Repeat  the  process  with  the  CaS04  and 
PbS04  in  the  other  tubes.     If,  however,  the  first  lot  fails  to  dissolve, 
ascertain  whether  any  has  dissolved  by  filtering  some  of  the  mixture 
very  carefully  and  evaporating  a  few  drops  of  the  filtrate  on  a  clean 
porcelain. 

148.  Solubility  in  hot  and  cold  water.     Weigh  four  portions  of  1  g. 
each  of  powdered  Ba(N03)2.     Place  10  cc.  of  water  in  a  test  tube  and 
add  one  portion  and  shake.     Note  the  result.     Warm  slowly,  and  as 
often  as  the  salt  is  entirely  dissolved  add  a  new  portion  of  1  g. 
Finally  bring  the  liquid  to  boiling. 

What  does  the  experiment  show  ? 

149.  Use  of  different  solvents.     Compare  the  solubility  of  iodine  in 
water,  carbon  disulphide,  and  alcohol.     Apply  no  heat.     Use  a  very 
small  quantity  of  I,  and  3  or  4  cc.  of  the  liquid.     Note  the  rapidity 
with  which  the  I  disappears,  and  judge  in  which  liquid  it  is  the  most 
soluble. 

150.  Compare  the  solubility  of  Nad  in  H20,  concentrated  HC1,  and 
alcohol. 

EXERCISE  XXXIX.     SOLUBILITY  (continued). 

151.  Solubility  in  mixtures.     Try  to  dissolve  BaCl2  in  concentrated 
HC1.     Finally  add  10  cc.  H20. 

152.  Dissolve  2  gr.  Na2S04  in  5  cc.  H20  and  add  an  equal  volume 
of  alcohol 

153.  Neutralization  of  the  solvents.  —  Dissolve  calcium  phosphate 
in  warm  dilute  HC1,  then  add  NH4OH  to  alkalinity.     Repeat,  using 
barium  oxalate. 

154.  Solution  of  liquids.      Test  the  solubility  of  the  following 
liquids  in  water :  alcohol,  ether,  olive  oil,  glycerol,  carbon  bisulphide. 


46  LABORATORY    EXERCISES. 

Proceed  in  each  case  as  follows :  Take  5  cc.  of  water  in  a  clean  test 
tube;  pour  1  cc.  of  the  liquid  to  be  tested  upon  the  water  in  the  tube. 
Shake  several  times,  and  then  observe  the  depth  of  the  liquid  layer 
above  or  below  the  water,  if  any  such  layer  there  be.  Ether  and  car- 
bon bisulphide  are  very  inflammable  and  very  volatile,  and  must  not  be 
brought  near  aflame. 

155.  Physical  and  chemical  solution.     Take  two  portions  of  5  gr. 
each  of  salsoda.     Dissolve  one  portion  in  10  cc.  of  water,  evaporate  to 
dryness  slowly,  and  compare  the  substance  obtained  with  the  original 
salt  in  appearance,  crystalline  form,  and  taste.     Dissolve  the  second 
portion  in  dilute  hydrochloric  acid,  evaporate,  and  compare. 

156.  Place  5  cc.  of  alcohol  in  a  test  tube  and  add  1  cc.  of  HC1,  and 
shake.    Drop  a  small  piece  of  fused  potassium  carbonate  into  the  tube. 
Place  5  cc.  of  water  in  a  test  tube  and  add  1  cc.  of  HC1.     Drop  a  small 
piece  of  fused  potassium  carbonate  into  this  tube.    Explain. 

EXERCISE  XL.     CRYSTALLIZATION. 

157.  Crystallization  by  solution  and  evaporation.     Place  5  gr.  of 
oxalic  acid  in  a  test  tube,  and  add  10  cc.  of  water.     Heat  until  the 
acid  has  dissolved,  and  then  allow  to  cool. 

158.  Crystallization  by  sublimation.     Heat  a  small  lump  of  benzoic 
acid  gently  in  a  dry  test  tube. 

159.  Crystallization  by  precipitation.     To  a  few  drops  of  concen- 
trated common  salt  solution  add  5  cc.  of  alcohol.     Examine  the  pre- 
cipitate under  a  microscope. 

160.  Purification  by  crystallization.     Weigh  25  gr.  of  soda  ash. 
Place  in  a  beaker,  and  pour  upon  the  ash  50  cc.  of  water.     Heat  until 
the  ash  has  dissolved,  filter  the  turbid  solution  while  hot,  and  allow  to 
cool.     Remove  some  of  the  crystals,  place  in  an  evaporating  dish,  and 
heat  until  they  are  reduced  to  a  fine  powder.     Compare  this  powder 
with  the  original  powder. 

LABORATORY  EXERCISES  ON  METALS. 

EXERCISE  XLI.    SODIUM  AND  POTASSIUM. 

161.  With  a  pair  of  pincers  remove  a  small  piece  of  metallic 
sodium,  Na,  from  the  oil,  and  dry  on  a  filter  paper.     Scrape  off  the 
outside  coating,  handling  the  metal  all  the  time  with  the  pincers. 
Note  whether  the  metal  is  hard  or  soft,  also  its  color  and  relative 


f  v- i  >    *  *  A?  T> 

f  or  THE    r   \ 

I   UNIVERSITY 


LABORATORY   EXERCISES.       V    P        OF  47 


weight  as  compared  with  water.  When  the  sodium  is  placed  on  water 
what  takes  place  ?  What  is  the  color  of  the  flame  ? 

When  action  ceases  test  the  solution  remaining  with  litmus  paper, 
and  draw  conclusion  as  to  class  of  compounds  to  which  it  belongs. 

Write  the  equation  for  the  action  of  sodium  on  water. 

162.  Perform  the  same  experiments  as  described  in  Exp.  161,  sub- 
stituting metallic  potassium,  K,  for  sodium.     Compare  results. 

Write  the  equation  for  the  action  of  K  on  H20,  for  the  combustion 
of  H,  and  for  the  combustion  of  K. 

163.  Test  some  potassium  carbonate,  K2C03,  solution  with  litmus 
paper.     Also  add  to  a  little  of  the  salt  a  few  drops  of  dilute  hydro- 
chloric acid.     Is  a  gas  given  off?    What  is  it? 

Treat  a  pound  of  wood  ashes  with  water,  and  filter  off  the  liquid, 
leaching  the  same  ashes  several  times  with  the  nitrate.  Examine  in 
the  same  manner  as  the  above  solution.  Evaporate  to  dryness,  and 
test  the  solid  remaining  with  a  little  dilute  hydrochloric  acid.  Does 
it  act  like  potassium  carbonate?  Save  the  material  for  use  in 
Exp.  165. 

164.  Examine  a  piece  of  caustic  potash  or  soda,  and  note  its  solu- 
bility.    Note  its  action  on  litmus  paper.     Place  the  solution  in  an 
evaporating  dish  and  add  slowly,  while  stirring  with  a  glass  rod, 
dilute  hydrochloric  acid  till  the  solution  is  slightly  acid.     Evaporate  to 
dryness.    What  is  the  powder  left  ?    Why  is  it  not  caustic  potash  ? 


EXERCISE  XLII 

165.  Tests  for  potassium.  Insert  a  piece  of  perfectly  clean  platinum 
wire  in  a  solution  of  potassium  chloride,  and  hold  it  in  the  non-lumi- 
nous flame  of  a  Bunsen  burner.  Note  the  color  imparted  to  the  flame. 

Try  the  same  with  a  sodium  chloride  solution.  Is  the  same  color 
imparted  to  the  flame  ? 

Mix  a  little  of  the  solutions  and  repeat  the  test.  Could  you  tell 
that  potassium  was  present  in  the  mixture  by  means  of  the  naked 
eye? 

Repeat  each  of  the  experiments,  using  blue  glass  through  which  to 
view  the  flame.  Describe  the  results. 

To  about  2  cc.  of  a  solution  of  potassium  chloride  add  a  like  quan- 
tity of  platinic  chloride.  What  is  the  result?  Try  the  same  with  a 
sodium  solution. 


48  LABORATORY   EXERCISES. 

From  the  above  experiments  explain  how  you  could  detect  the  dif- 
ference between  a  sodium  and  a  potassium  compound.  Try  the  tests 
on  a  solution  of  the  material  obtained  in  Exp.  163. 

EXERCISE  XLIII. 

166.  A  laboratory  use  of  borax.     What  is  the  formula  for  borax, 
and  what  is  its  other  chemical  name  ? 

Make  a  loop  in  a  piece  of  platinum  wire  as  large  as  a  capital  0  here 
printed.  Heat  the  loop  red  hot  in  the  flame,  and  thrust  it  while  still 
hot  into  some  powdered  borax,  a  quantity  of  which  will  adhere  to  the 
wire.  Reheat  the  loop  in  the  oxidizing  flame,  and  fuse  the  borax  to  a 
clear  glass.  When  a  good  bead  has  been  obtained  touch  it  while  still 
hot  to  a  very  small  particle  of  maganese  dioxide.  Reheat  in  the  oxidiz- 
ing flame  till  the  particle  is  seen  to  dissolve  and  diffuse  through  the 
bead.  Look  through  the  bead  toward  the  light.  What  is  the  color  ? 

Make  a  similar  experiment,  using  copper  oxide  instead  of  manganese 
dioxide. 

Ascertain  if  it  makes  any  difference  whether  the  oxidizing  or  the 
reducing  flame  be  used. 

For  what  purpose  might  borax  be  used  by  the  chemist  as  indicated 
by  the  above  experiment  ? 

EXERCISE  XLIV. 

167.  Name  the  principal  sodium  compound.     From  this  compound 
show  by  what  chemical  reactions  there  could  be  obtained  sodium  sul- 
phate, sodium  carbonate,  sodium  hydroxide. 

Start  with  20  g.  of  this  principal  salt,  and  using  such  other  material 
as  may  be  necessary,  make  some  potassium  chloride.  Prove  that  this 
is  a  potassium  compound. 

EXERCISE  XLV.     AMMONIUM. 

168.  Bring  near  to  each  other  two  vessels,  one  containing  a  little 
strong  ammonium  hydroxide  and  the  other  a  little  concentrated  hy- 
drochloric acid.     Explain  the  result. 

Try  strong  nitric  acid  instead  of  hydrochloric  acid.  Try  sulphuric 
acid.  How  do  you  explain  the  difference? 

169.  Re-read  Exp.  53  and  the  account  of  your  work  on  that  experi- 
ment.   How  could  ammonium  salts  be  detected? 


LABORATORY   EXERCISES.  49 

170.  In  a  dry  small  test  tube  heat  a  few  small  pieces  of  ammonium 
chloride.     What  collects  in  the  upper  part  of  the  tube  ?     Such  action 
is  called  sublimation.    Prove  by  some  test  that  it  is  an  ammonium  salt. 

Prove  that  it  is  ammonium  chloride  by  some  test.     (See  Exp.  61.) 

Ammonium  salts  of  volatile  acids  are  volatile. 

Ammonium  salts  of  non-volatile  acids  lose  their  ammonia  on  heating. 

All  ammonium  salts  are  decomposed  by  heating. 

171.  Dissolve  a  small  piece  of  alum  in  a  test  tube  and  add  am- 
monium hydroxide  till  after  shaking  the  solution  smells  slightly  of 
ammonia.     What  is  the  precipitate?    Write  the  equation.     What  is 
meant  by  a  precipitate  ?    Why  are  substances  thus  precipitated?    Judg- 
ing from  this  experiment,  for  what  purpose  could  ammonium  hydroxide 
be  used  in  the  laboratory  ? 

EXERCISE  XLVI.     CALCIUM. 

172.  Dissolve  about  10  gr.  of  marble  in  hydrochloric  acid.     Evapo- 
rate to  dryness,  and  allow  some  of  the  residue  to  be  exposed  to  the  air 
for  a  time.     Does  it  grow  moist?    What  is  the  compound  you  have 
made  and  which  thus  absorbs  the  moisture?    Write  the  equation  to 
show  the  reaction. 

If  you  should  add  concentrated  sulphuric  acid  to  some  of  this  resi- 
due, what  would  you  expect  to  be  the  result?  Try  it.  What  is  the 
solid  residue  from  the  sulphuric  acid  treatment?  Is  it  soluble  or 
insoluble  ? 

How  could  you  tell  whether  a  substance  is  potassium  chloride, 
sodium  chloride,  calcium  chloride,  or  ammonium  chloride. 

173.  Review  the  results  obtained  in  Exp.  120,  and  explain  how 
they  indicate  the  manner  of  cave  formation. 

174.  On  a  piece  of  quicklime  (what  is  the  formula?)  in  a  beaker 
pour  some  warm  water.     The  lime  slakes  and  falls  to  pieces.     Is  heat 
developed?      Does  the  lime  increase  in  bulk?      After  standing  for 
some  time  the  undissolved  lime  will  settle  out.     The  clear  liquid  is 
called  limewater. 

175.  Shake  a  little  powdered  plaster   of   Paris  (what  is  the  for- 
mula?) with  water  for  a  few  minutes.     Filter,  and  divide  the  nitrate 
into  four  parts,  using  one  part  for  each  of  the  following:  — 

(a)  To  one  add  a  few  drops  of  a  solution  of  barium  chloride.  A 
turbidity  of  barium  sulphate  shows  the  presence  of  sulphuric  acid. 

LAB.    EX.  — 4 


50  LABORATORY   EXERCISES. 

(&)  To  another  part  add  a  little  ammonium  oxalate  solution.  A 
white  precipitate  shows  lime.  The  two  tests  together  show  that  calcium 
and  sulphuric  acid  are  present.  If  both  are  present,  in  what  form  must 
they  be? 

(c)  To  a  third  portion  add  a  little  soap  solution.    A  scum  or  pre- 
cipitate indicates  a  hard  water. 

(d)  Boil  the  fourth  part.     No  precipitate  shows  permanent  hard- 
ness. 

What  is  the  cause  of  temporary  hardness  ?  (See  Exp.  120.)  How 
can  it  be  removed  ? 

EXERCISE  XLVH.     MAGNESIUM. 

176.  Characteristic  action  of  magnesium.     Pour  5  cc.  of  magnesium 
chloride  into  each  of  two  test  tubes.     To  one  add  a  little  ammonium 
hydroxide,  and  to  the  other  add  an  equal  volume  of  ammonium  chlo- 
ride and  then  a  little  ammonium  hydroxide. 

Repeat  the  experiment,  but  use  ammonium  carbonate  instead  of 
the  hydroxide.  What  is  the  precipitate  obtained  in  each  case  ?  Why 
is  a  precipitate  not  obtained  when  the  ammonium  compounds  are 
present? 

177.  To  3  cc.  of  a  magnesium  chloride  solution  add  about  5  cc. 
ammonium  chloride  and  an  equal  amount  of  ammonium  hydroxide. 
Dilute  with  an  equal  volume  of  water,  and  add  about  3  cc.  disodium 
phosphate,  Na2HP04,  solution  and  allow  to  stand.     A  crystalline  pre- 
cipitate thus  obtained  is  a  test  for  magnesium.     The  precipitate  is  am- 
monium-magnesium  phosphate.     Write    the    formula  for  it.     The 
precipitate  is  soluble  in  acids,  and  reprecipitated  by  ammonia. 

178.  Zinc.      Try  the  effect  of  nitric,  sulphuric,  and  hydrochloric 
acids,  dilute,  on  a  piece  of  metallic  zinc.     Describe  the  results. 

179.  In  a  solution  of  lead  acetate  (or  copper  sulphate)  immerse  a 
piece  of  sheet  zinc.     What  collects  on  the  zinc  ? 

180.  Place  5  cc.  zinc  sulphate  solution  in  a  test  tube  and  add  a  few 
drops  of  caustic  soda  solution;    finally  add  a  larger  quantity.     In 
excess  of  alkaline  hydroxides  soluble  zincates  are  formed. 

EXERCISE  XLVIII.     LEAD. 

181.  Examine  a  piece  of  sheet  lead.     Observe  its  color  and  softness. 
Heat  a  small  piece  on  charcoal  with  the  blowpipe.    Does  it  melt  easily  ? 
Try  its  solubility  in  nitric,  hydrochloric,  and  sulphuric  acids.     What 
is  the  compound  formed  in  each  case  ?    Is  each  soluble  ? 


LABORATORY   EXERCISES.  51 

182.  Heat  any  lead  compound  with  soda  on  charcoal.    What  do 
you  obtain  ?    How  do  you  know  ? 

183.  To  a  solution  of  a  lead  salt  add  dilute  hydrochloric  acid  till  no 
more  precipitate  forms.     Write  equation.     What  is  the  precipitate? 
Filter  off  the  precipitate,  and  to  one  third  of  nitrate  add  dilute  sul- 
phuric acid;  to  the  other  third  add  hydrogen  sulphide;  add  to  the  re- 
mainder potassium  chromate  solution.     Write  equation  for  each  test. 
Ascertain  if  the  original  precipitate  is  soluble  in  hot  water. 

If  the  precipitate  dissolves,  ascertain  if  soluble  chlorides  and  sul- 
phates coul,d  be  substituted  for  their  respective  acids  as  tests  for  lead. 

From  the  above  experiment  answer  the  following:  Is  lead  thrown 
out  of  a  solution  completely  by  hydrochloric  acid  or  a  soluble  chloride  ? 

How  could  lead  be  detected  in  an  unknown  solution  ? 

EXERCISE  XLIX.     MERCURY. 

184.  Examine  the  action  of  the  common  mineral  acids  on  mercury 
after  noting  the  physical  characteristics  of  the  metal.     Heat  a  little 
in  a  glass  tube  closed  at  one  end  and  note  the  result.     Drop  a  frag- 
ment of  zinc  into  a  little  mercury.    What  change  takes  place  ?    Define 
amalgam.     In  what  special  physical  property  does  this  metal  differ 
from  the  others  you  have  studied  ? 

185.  Put  into  each  of  two  test  tubes  4  or  5  cc.  of  water.    Add  to  one 
a  little  mercurous  chloride  and  to  the  other  mercuric  chloride,  and 
boil  each.     Which  dissolves? 

186.  In  one  of  two  test  tubes  place  about  5  cc.  of  mercuric  nitrate 
solution,  and  in  another  a  like  amount  of  mercurous  nitrate  solution, 
and  add  to  each  hydrochloric  acid  as  long  as  a  reaction  is  noted. 
Explain  the  result. 

In  the  same  way  try  the  action  of  these  salts  with  caustic  soda. 
What  is  formed  in  each  case  ? 

187.  Compare  the  action  of  the  two  salts  toward  hydrogen  sulphide. 

188.  To  a  solution  of  mercurous  nitrate  add  a  solution  of  stannous 
chloride.    What  is  the  precipitate  ?     Try  the  action  of  the  same  re- 
agent on  mercuric  chloride.     The  precipitate  is  at  first  white,  then 
gray.     This  is  a  characteristic  reaction. 

EXERCISE  L.     COPPER. 

189.  Examine  some  pieces  of  copper.     Try  their  solubility  in  the 
common  mineral  acids.    What  is  the  best  solvent? 


52  LABORATORY   EXERCISES. 

190.  With  the  blowpipe  heat  on  charcoal  a  little  copper  chloride 
with  an  equal  amount  of  sodium  carbonate.     Explain  the  result  in 
full. 

191.  To  a  solution  of  copper  chloride  add  sodium  carbonate  solu- 
tion so  long  as  a  precipitate  forms.     Boil  till  the  precipitation  turns 
black.     What  chemical  changes  take  place?      Filter,  dry,  and  save 
the  powder,  and  make  a  borax  bead  test  as  in  Exp.  166.     What  is  the 
color  of  the  bead  ?     This  is  a  characteristic  action  of  copper  compounds. 

192.  To  a  solution  of  copper  sulphate  add  some  caustic  potash 
solution.     What  is  the  result  ?    After  observing  the  characteristics  of 
the  precipitate,  heat,  and  notice  the  change.     Explain  the  action  and 
write  equations  for  the  reaction. 

193.  To  a  solution  of  copper  sulphate  add  ammonia  water,  drop 
by  drop,  shaking  the  solution  after  each  addition.     Note  all  changes. 
This  is  a  delicate  test  for  copper. 

194.  In  a  solution  of  copper  sulphate  dissolve  a  few  crystals  of  tar- 
taric  acid  and  then  add  an  excess  of  sodium  hydroxide.     Finally  add 
a  small  piece  of  glucose,  and  boil  the  solution.     What  takes  place  ? 
What  is  the  composition  of  the  precipitate  ?    What  are  the  formulae 
and  names  of  the  oxides  of  copper? 

195.  Acidify  5  cc.  of  copper  sulphate  solution  with  acetic  acid,  and 
add  a  few  drops  of  potassium  ferrocyanide  solution.     This  test  is  a 
more  delicate  one  for  copper  than  that  in  Exp.  193. 

196.  To  a  solution  of  copper  sulphate  add  a  little  hydrochloric 
acid.     Is  a  precipitate  formed  ?     Pass  sulphuretted  hydrogen  through 
the  solution.     What  is  the  result?     Filter   and  wash  thoroughly. 
Treat  the  precipitate  with  nitric  acid,  and  boil  with  water.     Test  for 
copper  as  indicated  in  the  previous  experiments. 

How  could  you  detect  copper  in  a  solution  ? 

EXERCISE  LI.    SILVER. 

197.  Fill  three  test  tubes  one  third  full  of  water  and  pour  into  each 
a  few  drops  of  silver  nitrate.     Add  to  one  about  3  cc.  of  hydrochloric 
acid,  and  shake  the  tube  violently.     To  the  second  tube  add  a  like 
quantity  of  potassium  bromide  solution  and  shake  the  tube.     To  the 
third^tube  add  a  like  quantity  of  potassium  iodide  solution.     Describe 
the  results  in  each  case  and  note  the  color  of  the  precipitates.     Name 
each  precipitate  and  write  the  equation  for  each  reaction. 

Withdraw  a  little  of  the  precipitate  from  each  tube  and  test  the 
solubility  in  nitric  acid. 


LABORATORY   EXERCISES.  53 

Test  also  the  solubility  of  the  precipitates  in  ammonium  hydroxide 
and  in  a  solution  of  sodium  thiosulphate. 

Heat  some  of  the  precipitate  from  one  tube  with  a  little  sodium 
carbonate  on  charcoal  before  the  blowpipe.  The  action  may  be  con- 
sidered typical  of  that  of  the  other  silver  compounds. 

198.  Precipitate  from  a  silver  solution  some  of  the  chloride  by 
means  of  hydrochloric  acid  (could  a  solution  of  sodium  chloride  be 
used  to  do  this?).  Throw  the  precipitate  on  a  filter  and  wash  it 
thoroughly  with  water.  Open  the  filter  and  spread  the  precipitate 
evenly  over  the  filter,  and  place  it  in  direct  sunlight.  Upon  the  facts 
illustrated  in  these  two  experiments  depends  the  art  of  photography. 


EXERCISE  LII.    ALUMINUM. 

199.  Examine  the  physical  properties  of  a  piece  of  aluminum.     Try 
the  action  of  the  common  mineral  acids  on  the  metal  as  in  the  .pre- 
ceding experiments  with  the  other  metals. 

200.  Dissolve  a  small  crystal  of  aluminum  sulphate  in  water,  in  a 
test  tube.    Place  one  half  of  the  solution  in  another  test  tube,  and  add 
drop  by  drop  ammonium  hydroxide,  till  after  shaking  the  odor  of 
ammonia  is  permanent.     Write  the  equation  for  the  reaction.     Place 
a  little  of  the  moist  hydroxide  in  another  tube  and  try  to  dissolve  it  in 
ammonium  hydroxide.     What  result?     In  a  similar  manner  try  its 
solubility  in  sodium  hydroxide.     Sodium  aluminate  is  formed. 

201.  With  the  remainder  of  the  solution  made  in  Exp.  200,  ascer- 
tain if  other  alkaline  hydroxides  will  precipitate  aluminum  hydroxide. 
Ignite  a  small  quantity  of  the  precipitate  on  charcoal  before  the  blow- 
pipe.    Moisten  the  solid  obtained  with  a  drop  of  cobalt  solution  ;  heat 
again  and  note  the  color  of  the  mass  obtained.     This  is  a  test  for 
aluminum. 

202.  Dissolve  a  little  "  alum  "  and  test  for  aluminum  as  indicated  in 
the  above  experiments. 

How  could  you  show  that  alum  contains  potassium  ?    That  it  is  a 
sulphate  ?    That  it  contains  water  of  crystallization  ? 

203.  To  a  small  quantity  of  an  organic  coloring  matter  (cochineal 
is  good),  add  a  little  potassium  sulphate  and  finally  a  little  ammonium 
hydroxide.     A  colored  precipitate  is  thrown  down,  which  is  known  as 
a  "  lake  "  color.     If  cochineal  is  used,  it  is  a  carmine  lake.    It  is 
simply  colored  aluminum  hydroxide. 


54  LABORATORY   EXERCISES. 

204.  Make  a  solution  of  aluminum  acetate  by  adding  to  a  solution 
of  lead  acetate  a  solution  of  alum  and  filtering  off  the  lead  sulphate. 
In  this  solution  soak  a  piece  of  cotton  cloth  and  put  it  away  till  the 
next  exercise. 

EXERCISE  LIE. 

204.  (Continued).     Treat  the  cotton  cloth  prepared  at  the  last 
exercise  as  well  as  a  piece  of  ordinary  cotton  cloth,  with  a  solution 
of  logwood,  and  observe  the  relative  amount  of  color  imparted.     To 
what  is  the  different  result  due  ?    What  is  a  material  that  will  act 
thus  called  ? 

IRON. 

205.  Dissolve  a  few  small  pieces  of  iron  (tacks  will  answer)  in  about 
10  cc.  of  dilute  sulphuric  acid.     When  action  ceases,  dilute  with  an 
equal  bulk  of  water.     What  is  given  off  during  the  action  ?    What  is 
in  solution  ?    Write  the  equation  for  the  action. 

Add  to  the  solution  ammonium  hydroxide  till  the  smell  of  am- 
monia is  permanent,  and  collect  the  precipitate  on  a  filter.  What  is 
the  precipitate?  What  is  its  color?  Write  the  equation  for  the 
reaction.  Ascertain  if  the  precipitate  is  soluble  in  hydrochloric  acid. 

206.  Ascertain  the  action  of  the  hydroxide  on  the  borax  bead  in  a 
loop  of  platinum  wire.    What  is  the  color  of  the  bead  ? 

207.  To  a  solution  of  ferrous  sulphate  add  ammonium  hydroxide  in 
excess.     Observe  color  of  precipitate.     Allow  it  to  stand  some  time, 
stirring  frequently,  and  observe  any  change  of  color.     Explain  the 
reaction. 

208.  Dip  a  piece  of  cotton  into  a  solution  of  nutgalls  or  tannin, 
and  dry  it.     Then  dip  it  into  a  solution  of  copperas  and  dry  again. 
Finally  try  to  wash  out  the  color. 

209.  Soak  a  piece  of  cotton  cloth  in  a  solution  of  ferric  sulphate 
and  then  immerse  it  in  an  acidulated  solution  of  potassium  ferro- 
cyanide.    Prussian  blue  is  precipitated  in  the  fibers  of  the  cloth. 


EXERCISE  LIV.    FERROUS  AND  FERRIC   COMPOUNDS. 

210.  Ferrous  salts.  Dissolve  a  little  iron  in  dilute  hydrochloric 
acid.  Place  half  the  solution  in  another  beaker  or  test  tube  for  use  in 
the  next  experiment.  To  a  part  of  the  solution  reserved  for  this 
experiment  add  a  few  drops  of  potassium  ferricyanide  solution ;  to 


LABORATORY   EXERCISES.  55 

another  part  add  a  few  drops  of  potassium  ferrocyanide  solution. 
Notice  the  difference.  To  a  third  portion  of  the  solution  add  a  few 
drops  of  potassium  sulphocyanide  solution. 

211.  Ferric  salts.     To  a  portion  of  the  solution  made  in  the  preced- 
ing experiment  and  reserved  for  this  experiment,  add  a  few  drops  of 
nitric  acid,  and  boil.     Test  this  solution,  as  in  the  preceding  experi- 
ment and  record  the  difference  in  action  as  compared  with  the  results 
obtained  above.     Explain  the  action  of  the  nitric  acid. 

212.  To  a  solution  of  a  ferric  salt  add  stannous  chloride  solution, 
and  warm.     Ascertain  whether  the  iron  is  in  the  ferric  or  the  ferrous 
condition.      How  may  ferrous  salts  be  changed  to  ferric  salts,  and 
vice  versa  ? 

In  what  ways  can  ferric  and  ferrous  salts  be  distinguished  from 
one  another  ?  What  is  the  most  delicate  test  for  each  class  of  these 
compounds  ? 

EXERCISE  LV.    ANALYSIS;  QUALITATIVE. 

Substances  may  be  subjected  to  analysis  with  two  ends  in  view. 

1.  To  ascertain  the  kind  of  matter  present  (Qualitative). 

2.  To  ascertain  the  quantity  of  each  kind  of  matter  (Quantita- 
tive). 

Review  the  experiments  recorded  in  your  notebook  and  make  a  list 
of  those  in  which  you  have  analyzed  qualitatively  any  substance. 
Do  the  same  with  experiments  in  which  quantitative  results  have  been 
obtained. 

In  qualitative  analysis  both  physical  and  chemical  changes  may  be 
used  as  tests. 

(a)  Physical  tests.  Such  tests  are  illustrated  in  Exp.  165.  Try  as 
there  described  for  sodium,  a  solution  of  a  lithium  salt,  also  of  a 
strontium  salt.  State  the  difference. 

(6)  Chemical  tests.  Such  tests  have  been  illustrated  several  times 
during  the  study  of  the  metals.  Look  over  your  previous  work  and 
select  at  least  three  such  cases.  These  may  be  further  illustrated 
by  the  following :  — 

213.  In   a  test  tube  place  about  5  cc.  of  silver  nitrate  solution 
and  a  like  quantity  of  copper  nitrate  solution.     Shake  well.     Proceed 
now  as  in  196.     What  is  the  precipitate  in  the  first  case?    What 
remains  in  the  filtrate  ?    What  is  the  precipitate  in  the  second  case  ? 
How  could  you  have  distinguished  a  lead  precipitate  from  one  of 
silver  in  the  first  instance  ?    (See  183.) 


56  LABORATORY   EXERCISES. 

214.  A  Qualitative  Analysis.    Put  a  ten-cent  silver  coin  into  an 
evaporating  dish,  and  pour  over  it  a  mixture  of  5  cc.  HN03  and  10  cc. 
H20.     Warm  it  till  all  or  nearly  all  has  dissolved. 

Remove  any  that  is  undissolved,  and  pour  the  liquid  into  a  test 
tube.  Add  HC1  as  long  as  a  precipitate  continues  to  form,  then  filter. 
AgCl  is  the  residue.  Give  the  reaction. 

Add  a  drop  or  two  of  HC1  to  the  filtrate,  and,  if  a  precipitate  falls, 
add  more,  and  filter  again,  to  remove  all  the  Ag. 

Evaporate  the  filtrate  to  a  few  drops  in  an  evaporating  dish  (to 
remove  any  free  HN03),  then  add  H20  and  pass  H2S  gas  into  the  filtrate 
so  long  as  a  precipitate  forms.  This  is  CuS.  Write  the  reaction. 
Filter.  The  coin  is  thus  found  to  contain  Ag  and  Cu. 

EXERCISE  LVI.    A  QUANTITATIVE  ANALYSIS. 

In  the  previous  experiments  what  substances  have  been  shown  to 
be  present  in  common  salt?  In  what  experiments  and  how  was  this 
shown  ?  By  what  kind  of  analysis  was  it  shown  ? 

215.  In  a  perfectly  dry  and  clean  mortar  pulverize  some  pure  so- 
dium chloride.     Put  it  in  an  evaporating  dish  and  heat  it  over  the 
Bunsen  flame  about  5  minutes.     Cool,  and  weight  out  accurately  on 
the  balance  1  g.  of  the  salt.   Dissolve  it  in  a  beaker  by  means  of  about 
25  cc.  of  distilled  water.    Warm  the  solution  and  add  as  long  as  a 
precipitate  is  formed  a  silver  nitrate  solution,  pouring  it  gradually 
and  with  constant  stirring. 

The  contents  of  the  beaker  should  be  protected  from  the  direct  sun 
light  as  much  as  possible.  Allow  the  precipitate  to  settle,  and  decant 
the  supernatant  liquid  through  a  well-dried  and  weighed  filter  held  in 
a  funnel.  Care  should  be  taken  to  make  sure  that  all  the  chlorine  has 
been  precipitated.  Finally,  transfer  the  whole  of  the  precipitate  to  the 
filter  paper,  being  careful  not  to  lose  even  the  least  particle.  Wash 
the  precipitate  on  the  filter  thoroughly.  Dry  on  the  filter  paper  to 
constant  weight,  being  careful  not  to  scorch  the  paper.  The  increase 
in  weight  is  the  weight  of  the  silver  chloride.  Calculate  now  the  per 
cent  of  chlorine  in  silver  chloride,  AgN03,  and,  using  this  as  a  factor, 
calculate  the  per  cent  of  chlorine  in  the  sodium  chloride  used,  and 
also  the  per  cent  of  sodium. 


APPENDIX. 

THE  METRIC   SYSTEM  OF  WEIGHTS  AND  MEASURES.1 

The  metric  system,  employed  in  the  affairs  of  everyday  life  by 
most  of  the  nations  of  continental  Europe,  and  by  scientific  writers 
throughout  the  world,  is  based  upon  a  fundamental  unit,  or  measure 
of  length,  called  a  meter.  This  meter  is  defined  as  the  40-millionth 
part  of  the  circumference  of  the  earth,  or,  in  other  words,  of  a  "  great 
circle,"  or  meridian.  Its  length  was  originally  determined  by  actual 
measurement  of  a  considerable  arc  of  a  meridian,  but  the  various 
measurements  heretofore  made  of  the  length  of  the  earth's  meridian 
differ  slightly  from  each  other ;  and  it  is  to  be  expected,  and  indeed 
hoped,  that  the  steady  improvements  of  methods  and  instruments 
will  make  each  successive  determination  of  the  length  of  the  meridian 
better  than,  and  therefore  different  from,  the  preceding.  It  is  on  this 
account  necessary  to  define  the  standard  of  length  by  legislation  to 
be  a  certain  rod  of  metal,  deposited  in  a  certain  place,  under  specified 
guaranties,  and  to  secure  the  uniformity  and  permanence  of  the 
standard  by  the  multiplication  of  exact  copies  in  safe  places  of 
deposit. 

From  this  single  quantity,  the  meter,  all  other  measures  are  deci- 
mally derived.  Multiplied  or  divided  by  10,  100,  1000,  and  so  forth, 
the  meter  supplies  all  needed  linear  measures ;  and  the  square  meter 
and  cubic  meter,  with  their  decimal  multiples,  supply  all  needed 
measures  of  area  or  surface  on  the  one  hand,  and  of  solidity  or 
capacity  on  the  other. 

From  the  unit  of  measure  to  the  unit  of  weight,  the  transition  is 
admirably  simple  and  convenient.  The  cube  of  the  one  hundredth  of 
the  linear  meter  is,  of  course,  the  millionth  of  the  cubic  meter :  its 
bulk  is  about  that  of  a  large  die  of  the  common  backgammon  board. 
This  little  cube  of  pure  water  is  the  universal  unit  of  weight,  a  gram, 
which,  decimally  multiplied  and  divided,  is  made  to  express  all 
weights.  The  numbers  expressing  all  weights,  from  the  least  to  the 
greatest,  find  direct  expression  in  the  decimal  notation ;  the  weights 
used  in  different  trades  only  differ  from  each  other  in  being  different 

iFrom  Manual  of  Chemistry,  by  Storer  and  Lindsay. 
67 


58 


LABORATORY   EXERCISES. 


decimal  multiples  of  the  same  fundamental  unit ;  and,  in  comparing 
together  weights  and  volumes,  none  but  easy  decimal  computations 
are  ever  necessary. 

The  nomenclature  of  the  metrical  system  is  extremely  simple ;  one 
general  principle  applies  to  each  of  the  following  tables.  The  Greek 
prefixes  for  10,  100,  and  1000,  viz.,  deca,  hecto,  and  kilo,  are  used  to 
signify  multiplication;  while  the  Latin  prefixes  for  10,  100,  and  1000, 
viz.,  deci,  centi,  and  milli,  are  employed  to  express  subdivision.  Of 
the  names  thus  systematically  derived  from  that  of  the  unit  in  each 
table,  many  are  not  often  used ;  the  names  in  common  use  are  those 
printed  in  small  capitals.  Thus,  in  the  table  for  linear  measure,  only 
the  meter,  kilometer,  centimeter,  and  millimeter  are  in  common  use, 
—  the  first  for  such  purposes  as  the  English  yard  subserves,  the 
second  instead  of  the  English  mile,  the  third  and  fourth  in  lieu  of  the 
fractions  of  the  English  foot  and  inch. 


Divisions 
Unit  .  . 
Multiples 

Divisions 
Unit  .  . 

Divisions 
Unit  .  . 
Multiples 


LINEAR  MEASURE. 

Meters. 

f  MILLIMETER 

0.001  or  1-1000 

of  a  meter. 

j  CENTIMETER 

0.01    or  1-100 

« 

[  Decimeter 

0.1      or  1-10 

« 

METER 

1. 

f  Decameter 

10. 

\  Hectometer 

=       100. 

I  KILOMETER 

=     1000. 

SURFACE  MEASURE. 

f  Millimeter  square  =          0.000,001  of  a  meter  square. 
|  Centimeter  square  =          0.000,1  "  " 

I  Decimeter  square  =          0.01  "  " 

METER  SQUARE    =          1. 


f  Cubic  Millimeter  = 
j  Cubic-  Centimeter  = 
[  Cubic  Decimeter  = 
CUBIC  METER  = 
C  Cubic  Decameter  = 
I  Cubic  Hectometer  = 


CUBIC  MEASURE. 

Cubic  Meters. 

0.000,000,001 

0.000,001 
0.001 

1. 

1,000. 
1,000,000. 


i  Cubic  Kilometer    =1,000,000,000. 


APPENDIX. 


59 


The  table  for  land  measure  we  omit,  as  having  no  connection  with 
our  subject.  For  the  measurement  of  wine,  beer,  oil,  grain,  and  simi- 
lar wet  and  dry  substances,  a  smaller  unit  than  the  cubic  meter  is 
desirable.  The  cubic  decimeter  has  been  selected  as  a  special  stand- 
ard of  capacity  for  the  measurement  of  substances,  such  as  are  bought 
and  sold  by  the  English  wet  and  dry  measures.  The  cubic  decimeter 
(1000  cubic  centimeters)  thus  used  is  called  a  liter. 

CAPACITY  MEASURES. 

Liters.  Cubic  Meter. 

f  Milliliter  =       0.001  =  0.000,001  =  1  cubic  centimeter. 

Divisions  .  |  Centiliter  =        0.01  =  0.000,01 

[Deciliter  =        0.1  =0.000,1 

Unit     .    .     LITER  =        1.  =  0.001        =  1  cubic-decimeter. 

f  Decaliter  =  10.  =0.01 

Multiples  .  ]  HECTOLITER  =  100.  =0.1 

iKiloliter  =1,000.  =1. 


=  1  cubic  meter. 


The  table  of  weights  bears  an  intimate  relation  to  this  table  of 
capacity.  As  already  mentioned,  the  weight  of  that  die-sized  cube,  a 
cubic  centimeter,  or  milliliter,  of  distilled  water  (taken  at  4°,  its  point 
of  greatest  density),  constitutes  the  metric  unit  of  weight.  This 
weight  is  called  a  gram.  From  the  very  definition  of  the  gram, 
and  from  the  table  of  capacity  measure,  it  is  clear  that  a  liter  of  dis- 
tilled water  at  4°  will  weigh  1,000  grams. 

WEIGHTS. 

Grams. 

f  MILLIGRAM  =  0.001 

Divisions  .  |  CENTIGRAM  =  0.01 

[  DECIGRAM    =  0.1 

Unit     .    .     GRAM           =  1.  =  1  cubic  centimeter  of  water  at  4°. 

f  Decagram     =  10. 

Multiples  .  I  Hectogram  =  100. 

I  Kilogram     =  1,000.  =  1  cubic  decimeter  of  water  at  4°. 

The  simplicity  and  directness  of  the  relations  between  weights  and 
volumes  in  the  metric  system  can  now  be  more  fully  explained.  The 
chemist  ordinarily  uses  the  gram  as  his  unit  weight,  and,  for  his 
unit  of  volume,  a  cubic  centimeter,  which  is  the  bulk  of  a  gram 


60  LABORATORY    EXERCISES. 

of  water.  For  coarser  work,  the  kilogram  becomes  the  unit  of 
weight,  and  the  corresponding  unit  of  measure  is  the  liter,  which  is 
the  bulk  of  a  kilogram  of  water.  In  commercial  dealings,  in  manu- 
facturing processes,  and,  above  all,  in  scientific  investigations,  these 
simple  relations  between  weights  and  measures  have  been  found  to 
be  an  inestimable  advantage.  The  numerical  expressions  for  metric 
weights  and  measures  may  always  be  read  as  decimals.  Thus,  5.126 
meters  will  be  read,  "  five  meters  and  one  hundred  and  twenty-six 
thousandths,"  and  not  "  five  meters,  one  decimeter,  two  centimeters, 
and  six  millimeters."  The  expression  "  10.5  grams  "  is  read  "  ten  and 
five  tenths  grams  " ;  just  as  we  say  one  hundred  and  five  dollars,  not 
ten  eagles  and  five  dollars;  or  sixty-five  cents,  not  six  dimes  and  five 
cents.  All  computations  under  the  metric  system  are  made  with 
decimals  alone. 

The  abbreviations  commonly  met  with  in  chemical  literature  are :  — 

mm.  for  millimeter.  cm.  for  centimeter, 

m.  for  meter.  cc.  or  cm8  for  cubic  centimeter, 

g.  or  grm.  for  gram.  k.  or  kgm.  or  kilo,  for  kilogram. 
1.  for  liter. 

The  equivalents  in  English  weights  and  measures,  of  those  metric 
weights  and  measures  which  are  used  in  chemistry,  can  be  readily 
found  by  the  aid  of  the  table  on  page  63,  which  is  available 
not  only  for  grams,  centimeters,  and  liters,  but,  by  mere  change  of 
the  position  of  the  decimal  point,  for  all  decimal  multiples  or  subdi- 
visions of  these  quantities. 

One  cubic  meter  =  35.31660  cubic  feet. 

One  cubic  decimeter  (a  liter)  =  61.02709  cubic  inches. 

One  cubic  centimeter  =        0.06103  cubic  inch. 

One  liter  =         0.22017  imperial  gallon. 

One  liter  =         0.88066  imperial  quart. 

One  liter  =         1.76133  imperial  pint. 

One  liter  =        0.26427  U.  S.  gallon. 

One  liter  =        1.05708  U.  S.  quart. 

One  liter  =        2.11415  U.  S.  pints. 

One  gram  =  15.4323  grains. 

One  meter  =  39.3708  inches. 


APPENDIX. 


61 


One  pound  Avoirdupois          =  7,000         grains 
One  pound  Troy  =  5,760        grains 

One  ounce  Avoirdupois  =     437.5     grains 

One  ounce  Troy  =     480         grains 

One  grain 

One  English  imperial  gallon  =     277.274  cu.  in. 
One  U.  S.  standard  gallon      =     231        cu.  in. 
One  U.  S.  quart 
One  fluid  ounce 
One  foot 
One  yard 
One  inch 


453.59  g. 

373.24  g. 

28.35  g. 

31.10  g. 

64.80  mg. 

4.54  1. 

3.78  1. 

0.95  1. 
29.56  cc. 

0.3048  m. 

0.9144  m. 

2.54  cm. 


LABORATORY  EXERCISES. 


TABLE 

FOB   THE   CONVERSION   OF   DEGREES    ON   THE    CENTIGRADE    THERMOMETER 
INTO   DEGREES   OF   FAHRENHEIT'S    SCALE. 

C.°  =  (F.°  -  32)  f ;  and  F.°  =  f  C.°  +  32. 


CENT. 

FAHR. 

CENT. 

FAHR. 

CENT. 

FAHR. 

o 
-50 

0 

-58.0 

0 

17 

0 

62.6 

o 
60 

140.0 

-45 

-49.0 

18 

64.4 

61 

141.8 

-40 

-40.0 

19 

66.2 

62 

143.6 

-35 

-31.0 

20 

68.0 

63 

146.4 

-30 

-22.0 

21 

69.8 

64 

147.2 

-25 

-13.0 

22 

71.6 

65 

149.0 

-20 

-  4.0 

23 

73.4 

66 

150.8 

-19 

-  2.2 

24 

75.2 

67 

152.6 

-18 

-  0.4 

25 

77.0 

68 

154.4 

-17 

+  1.4 

26 

78.8 

69 

156.2 

-16 

3.2 

27 

80.6 

70 

158.0 

-15 

5.0 

28 

82.4 

71 

159.8 

-14 

6.8 

29 

84.2 

72 

161.6 

-13 

8.6 

30 

86.0 

73 

163.4 

-12 

10.4 

31 

87.8 

74 

165.2 

-11 

12.2 

32 

89.6 

75 

167.0 

-10 

14.0 

33 

91.4 

76 

168.8 

-  9 

15.8 

34 

93.2 

77 

170.6 

-  8 

17.6 

35 

95.0 

78 

172.4 

-  7 

19.4 

36 

96.8 

79 

174.2 

-  6 

21.2 

37 

98.6 

80 

176.0 

-  5 

23.0 

38 

100.4 

81 

177.8 

-  4 

24.8 

39 

102.2 

82 

179.6 

-  3 

26.6 

40 

104.0 

83 

181.4 

-  2 

28.4 

41 

105.8 

84 

183.2 

-  1 

30.2 

42 

•  107.6 

85 

185.0 

0 

32.0 

43 

109.4 

86 

186.8 

+  1 

33.8 

44 

111.2 

87 

188.6 

2 

35.6 

46 

113.0 

88 

190.4 

3 

37.4 

46 

114.8 

89 

192.2 

4 

39.2 

47 

116.6 

90 

194.0 

5 

41.0 

48 

118.4 

91 

195.8 

6 

42.8 

49 

120.2 

92 

197.6 

7 

44.6 

60 

122.0 

93 

199.4 

8 

46.4 

61 

123.8 

94 

201.2 

9 

48.2 

62 

125.6 

95 

203.0 

10 

60.0 

63 

127.4 

96 

204.8 

11 

51.8 

64 

129.2 

97 

206.6 

12 

53.6 

55 

131.0 

98 

208.4 

13 

55.4 

66 

132.8 

99 

210.2 

14 

57.2 

67 

134.6 

100 

212.0 

15 

69.0 

58 

136.4 

16 

60.8 

69 

138.2 

APPENDIX. 


63 


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a 

STORER    AND    LINDSAY'S 

Elementary  Manual  of  Chemistry 

BY  F.  H.  STORER,  S.B.,  A.M.,  and  W.  B.  LINDSAY,  A.B.,  B.S. 
Cloth,  12mo,  453  pages.     Illustrated.     Price,  $1.20 

This  work  is  the  lineal  descendant  of  the  "  Manual  of 
Inorganic  Chemistry"  of  Eliot  and  Storer,  and  the  "Ele- 
mentary Manual  of  Chemistry  "  of  Eliot,  Storer  and  Nichols. 
It  is  in  fact  the  last  named  book  thoroughly  revised, 
rewritten  and  enlarged  to  represent  the  present  condition 
of  chemical  knowledge  and  to  meet  the  demands  of  American 
teachers  for  a  class  book  on  Chemistry,  at  once  scientific 
in  statement  and  clear  in  method. 

The  purpose  of  the  book  is  to  facilitate  the  study  and 
teaching  of  Chemistry  by  the  experimental  and  inductive 
method.  It  presents  the  leading  facts  and  theories  of  the 
science  in  such  simple  and  concise  manner  that  they  can 
be  readily  understood  and  applied  by  the  student.  The 
book  is  equally  valuable  in  the  classroom  and  the  laboratory. 
The  instructor  will  find  in  it  the  essentials  of  chemical 
science  developed  in  easy  and  appropriate  sequence,  its 
facts  and  generalizations  expressed  accurately  and  scientifi- 
cally as  well  as  clearly,  forcibly  and  elegantly. 


' '  It  is  safe  to  say  that  no  text-book 
has  exerted  so  wide  an  influence 
on  the  study  of  chemistry  in  this 
country  as  this  work,  originally 
written  by  Eliot  and  Storer.  Its 
distinguished  authors  were  leaders 
in  teaching  Chemistry  as  a  means 
of  mental  training  in  general  edu- 
cation, and  in  organizing  and  per- 
fecting a  system  of  instructing 
students  in  large  classes  by  the 
experimental  method.  As  revised 
and  improved  by  Professor  Nichols, 
it  continued  to  give  the  highest 
satisfaction  in  our  best  schools  and 
colleges.  After  the  death  of  Pro- 
fessor Nichols,  when  it  became 


necessary  to  revise  the  work  again, 
Professor  Lindsay,  of  Dickinson 
College,  was  selected  to  assist  Dr. 
Storer  in  the  work.  The  present 
edition  has  been  entirely  rewritten 
by  them,  following  throughout  the 
same  plan  and  arrangement  of  the 
previous  editions,  which  have  been 
so  highly  approved  by  a  generation 
of  scholars  and  teachers. 

"  If  a  book,  like  an  individual, 
has  a  history,  certainly  the  record 
of  this  one,  covering  a  period  of 
nearly  thirty  years,  is  of  the  highest 
and  most  honorable  character." 
— From  The  American  Journal  of 
Science. 


Copies  cf  this  book  will  be  sent  prepaid  to  any  address^  on  receipt  of  the  price^ 
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